:-NRLF 


E77    fl?D 


INTERNATIONAL  ATOMIC   WEIGHTS,    1915 


Symbol. 

Atomic 
weight. 

Symbol. 

Atomic 
weight. 

Aluminum 
Antimony 

Al 

27.1 
120.2 
39.88 
74.96 
137.37 
9,1 
208.0 
11.0 
79.92 
112.40 
132.81 
40.07 
12.00 
140.25 
35.46 
52.0 
58.97 
93.5 
63.57 
162.5 
167.7 
152.0 
19.0 
157.3 
69.9 
72.5 
197.2 
3.99 
163.5 
1.008 
114.8 
126.92 
193.1 
55.84 
82.92 
139.0 
207.10 
6.94 
174.0 
24.32 
54.93 
200.6 

Molybdenum   .....  Mo 
Neodymium     ...             Nd 
Neon  Me 
Nickel    Ni 
Nitoa  (radium  emanation)  Nt 
Nitrogen   N 

96.0 
144.3 
20.2 
58.68 
222.4 
14.01 
190.9 
16.00 
106.7 
31.04 
195.2 
39.10 
140.6 
226.4 
102.9 
85.45 
101.7 
150.4 
44.1 
79.2 
28.3 
107.88 
23.00 
87.63 
32.07 
181.5 
127.5 
159.2 
204.0 
232.4 
168.5 
119.0 
48.1 
184.0 
238.5 
51.0 
130.2 

172.0 
89.0 
65.37 
90.6 

Sb 
A 

Arsenic   .    . 

As 

Ba 

Be 

Bi 

.  .       Os 

B 

O 

Br 

Palladium  

.  .  .   Pd 

.  .           Cd 

Phosphorus      .  .  . 

.  .   .   P 
.  .       Pt 

Caesium 

Cs 

Calcium 

Ca 

.  .   .  K 

Carbon 

c 

Praseodymium    .   . 
Radium    
Rhodium     .... 

.  .  .    Pr 
.   .       Ra 
.  .  .   Rh 
.   .      Rb 

Ce 

Chlorine 

Cl 

Cr 

Cobalt 

Co 

.   .  .  Ru 

Columbium 

Cb 

.      .  Sa 

Cu 

Scandium     .... 

.  .   .  Sc 

Dysprosium 

Dv 

.    .   .   Se 
.    .  .  Si 

Er 

Eu 

Silver     .  . 

Ac 

F 

.  .  .  Na 

Gadolinium 

Gd 

.   .   .  Sr 

Ga 

Sulphur     

.   .  .  S 

Germanium 
Gold 

Ge 

Tantalum     .... 

Ta 

Au 

Tellurium    .... 

.   .   .  Te 

He 

.  .   .  Tb 

Holmium    . 

Ho 

Thallium  .      ... 

.   .   .  Tl 
.       .  Th 
.   .  Tm 

H 

Indium 

...             In 

I 

Tin 

Sn 

Ir 

Ti 

Fe 

W 

Krypton  .   . 
Lanthanum 
Lead     .   .  . 
Lithium  .   . 

Kr 

La 

Uranium  

.  .   .  U 
.   .   .  V 

Pb 
Li 

.   .       Xe 

Ytterbium 
(Neoytterbium) 
Yttrium  .   ... 
Zinc          

.  .   .   Yb 
.   .  .  Yt 
.   .       Zn 

Lutecium    . 

Lu 

Magnesium 
Manganese 

Mg 

Mn 

He 

Zirconium    .... 

.   .  .  Zr 

From  the  collection  of  the 


n 


PreTinger 
v    iJibrary 


San  Francisco,  California 
2006 


SYNTHETIC 
INORGANIC  CHEMISTRY 


A    LABORATORY    COURSE   FOR 
FIRST   YEAR   COLLEGE    STUDENTS 


BY 


ARTHUR   A.    BLANCHARD,    PH.D. 

A  ssociate  Professor  of  Inorganic  Chemistry  at  the 
Massachusetts  Institute  of  Technology 


SECOND    EDITION,    WITH   SUPPLEMENT 
FOURTH  THOUSAND 


NEW  YORK 

JOHN  WILEY  &  SONS,  INC. 

LONDON  :  CHAPMAN  &  HALL,  LIMITED 

1916 


COPYRIGHT,  1908,  1910,  1916 

BY 
ARTHUR    A.    BLANCHARD 


Stanbope  ]press 

F.    H.   GILSON     COMPANY 
BOSTON.     U.S.A. 


PREFACE 

THIS  series  of  notes  was  designed  to  serve  as  a  guide 
for  laboratory  work  and  study  in  Inorganic  Chemistry  during 
the  second  term  of  the  first  year  at  the  Massachusetts  Insti- 
tute of  Technology.  It  had  been  felt  for  some  time  that 
Qualitative  Analysis,  which  was  previously  made  the  basis 
for  laboratory  practice  during  that  period,  did  not  fully 
meet  the  requirements  and  that  a  course  based  upon  the 
actual  preparation  of  typical  chemical  substances  might 
prove  more  satisfactory.  In  consequence,  notes  in  essen- 
tially the  form  now  published  were  prepared  during  the  year 
1906—07,  they  being  the  direct  outcome  of  several  years'  pre- 
vious trial  of  a  limited  amount  of  preparation  work  with  the 
classes.  The  present  book  is  a  thorough  revision  of  those 
notes  in  the  light  of  experience  in  their  actual  application. 

During  the  first  term's  study  of  chemistry  there  can  be 
little  doubt  that  a  course  of  simple  experiments,  such  as  has 
long  been  in  use,  in  the  methods  of  formation  and  in  the 
study  of  the  properties  of  the  non-metallic  elements  —  oxy- 
gen, hydrogen,  the  halogens,  sulphur,  nitrogen,  and  carbon 
—  and  their  compounds,  is  the  most  satisfactory.  But  when 
it  comes  to  the  study  of  the  metallic  elements,  three  options 
as  to  laboratory  work  present  themselves :  First,  a  continua- 
tion of  experiments  similar  in  nature  to  those  of  the  first  term  ; 
second,  Qualitative  Analysis ;  third,  Preparation  Work.  The 
disadvantages  of  the  first  plan  are  that  the  experiments  are 
so  quickly  performed  and  so  alike  in  character  that  they 
fail  to  arouse  much  enthusiasm  in  the  student  or  to  leave 
very  vivid  impressions  on  his  mind.  .Qualitative  analysis 
is  in  many  ways  a  most  excellent  basis  for  teaching  the 
chemistry  of  the  metallic  elements ;  but  its  chief  disadvan- 
tages are :  First,  that  it  is  one-sided,  it  dealing  as  it  does 

iii 

304847 


IV  PREFACE 

almost  exclusively  with  the  chemistry  of  solutions  and  the 
formation  of  highly  insoluble  bodies ;  second,  that  it  requires 
the  sequence  followed  in  the  lectures  to  be  that  of  the  quali- 
tative procedure  instead  of  a  more  natural  one  based  on  the 
periodic  classification ;  and  third,  that  it  is  well-nigh  impos- 
sible to  keep  from  the  student's  mind  the  false  idea  that  the 
end  and  aim  of  qualitative  analysis  is  principally  "  to  get 
the  unknowns  correct." 

Some  of  the  advantages  which  seem  to  be  possessed 
by  a  course  of  preparation  work  such  as  outlined  in  the 
following  pages  are: 

1.  The  sequence   of  the  exercises  may  follow  that  of 
the  lectures. 

2.  Very  varied  types  of  chemical  change  are  illustrated, 
both  those  in  the  furnace  and  those  in  solution.     In  solu- 
tion advantage  is  taken  not  only  of  high  degrees  of  insol- 
ubility, but  also  of  differences  in  solubility  among  the  more 
soluble   bodies  as  well   as  of   differences  in   the  effect   of 
temperature  on  solubility. 

3.  The    danger   of   the   work   becoming   a    mechanical 
following  of   directions    is  reduced   by  the  introduction  of 
study  questions  and  experiments  with  each  exercise. 

4.  In  its  effect  in  awakening  the  student's  interest  this 
line  of  work  has  proved  particularly  successful,  —  the  making 
of  preparations  is,  in  fact,  in  its  very  nature  one  of  the  most 
fascinating  forms  of  chemical  work.     Since  each  preparation 
requires  a  good  deal  of  time  and  thought,  and  the  product 
when  obtained  is  something  definite  and  tangible,  the  knowl- 
edge thus  gradually  absorbed  is  more  definite  and  less  easily 
forgotten  than  when  the  laboratory  work  consists  of  a  large 
number  of  test  tube  reactions. 

After  the  completion  of  such  a  course  as  this,  if  the 
student  commences  analytical  work  with  some  conception 
of  the  sources  and  methods  of  obtaining  the  substances 


PREFACE  V 

which  he  is  to  use  as  reagents,  etc.,  there  can  be  no  doubt 
that  the  latter  work  will  then  have  a  much  deeper  meaning. 

The  plan  kept  in  mind  in  preparing  this  course  is,  briefly, 
as  follows :  The  greater  part  of  the  preparations  selected  are 
of  industrial  importance,  and  for  the  starting  point  of  each 
either  natural  products  or  crude  manufactured  materials  are 
used  so  far  as  is  possible.  The  course  does  not  aim  to  be 
an  exhaustive  one  in  chemical  preparations,  but  a  limited 
number  of  exercises  are  selected  to  illustrate  the  most  impor- 
tant types  of  compounds  of  the  common  elements  and  the 
most  important  methods.  Two  Or  three  times  as  many 
exercises  are  furnished  as  any  one  student  will  be  able  to 
complete  in  the  time  usually  allotted ;  thus  different  students 
may  be  assigned  different  preparations. 

The  notes  for  each  exercise  are  divided  into  three  parts : 
I.  A  discussion  of  the  object  of  the  exercise,  with  an  out- 
line of  the  principle  of  the  method  and  the  reasons  for  the 
steps  involved.  II.  Working  directions  which,  if  carefully 
followed,  should  result  in  obtaining  a  satisfactory  product. 
It  is  believed  far  better  to  make  the  directions  very  explicit, 
for  the  reason  that  the  inexperienced  student  may  easily 
become  discouraged  by  failures  due  to  difficulties  which  he 
is  unable  to  foresee.  Difficulties  enough  are  sure  to  arise 
to  develop  originality  and  resourcefulness.  III.  Questions 
for  study  which  involve  additional  laboratory  experiments, 
the  consulting  of  text-books,  and  original  reasoning. 

At  the  end  of  each  group  of  exercises  is  furnished  a  set 
of  general  study  questions,  and  this  arrangement  of  the 
exercises  in  groups  is  such  as  to  bring  out  the  relationships 
shown  in  the  periodic  classification  of  the  elements. 

In  the  discussions  and  questions  given  with  the  various 
exercises  it  is  assumed  that  the  student  has  an  elementary 
knowledge  of  the  electrolytic  dissociation  theory  and  of 
the  principle  of  mass  action.  In  the  opinion  of  the  author 


Vl  PREFACE 

a  great  opportunity  is  lost  for  bringing  out  relationships 
among  chemical  phenomena  if  these  principles  are  not 
taught  during  the  first  term's  study  of  college  chemistry 
and  their  applications  pointed  out  in  connection  with  later 
work  in  inorganic  preparations  and  in  analytical  chemistry. 
The  effort  has  been  to  make  the  questions  such  as  cannot 
be  answered  mechanically.  Some  of  the  questions  may,  in 
consequence,  seem  rather  difficult  and  incapable  of  direct 
answers ;  the  object  of  the  questions  is,  however,  not  solely 
to  bring  forth  correct  statements  of  facts  and  theories,  but 
is  also  to  teach  the  student  to  use  his  head  in  seeking  for 
the  significance  of  facts  and  in  reasoning  from  one  fact  to 
another. 

Acknowledgment  is  due  to  many  sources  for  the  outline 
of  the  greater  part  of  the  methods  given.  The  details  of 
all  of  them  have,  however,  been  very  carefully  worked  over 
and  adapted  for  the  purpose  in  view. 

In  conclusion  the  author  wishes  to  express  his  sense  of 
obligation  to  Professor  Henry  P.  Talbot,  head  of  the  Depart- 
ment of  Chemistry,  at  whose  request  the  preparation  of 
these  notes  was  undertaken ;  also  to  other  members  of  the 
instructing  staff  at  the  Massachusetts  Institute  of  Tech- 
nology for  helpful  criticism  and  suggestions,  and  particularly 
to  Mr.  J.  W.  Phelan,  to  whose  efficient  management  of  the 
laboratory  instruction  is  due  any  success  with  which  this 
course  has  met  at  this  Institute. 

This  little  book  is  presented  for  publication  with  the 
desire  to  offer  for  the  consideration  of  those  in  charge  of 
the  instruction  at  other  institutions  a  plan  of  first-year  work 
which  has  quite  new  and  perhaps  advantageous  features. 
It  is  hoped  that  it  may  do  its  part  in  securing  recognition 
of  the  importance  of  synthetic  or  preparative  work  in  a 
well-balanced  course  of  chemical  training. 

ARTHUR  A.  BLANCHARD. 

March,  1908. 


PREFACE   TO    SECOND    EDITION 

THE  plan  of  work  embodied  in  the  first  edition  has 
remained  unchanged  in  the  second  edition,  but  many  minor 
improvements  have  been  made  and  considerable  new  mate- 
rial has  been  added. 

Several  of  the  procedures  have  been  altered,  so  that 
good  results  may  be  more  confidently  expected  from  inex- 
perienced students.  A  few  new  preparations  have  been 
added  to  the  list  in  order  to  include  important  types  of 
processes  which  were  not  well  represented  in  the  first  edi- 
tion. The  general  questions  following  each  group  of  exer- 
cises have  been  entirely  rewritten,  and  now  present  a  more 
consistent  plan  to  bring  out  the  main  characteristics  of  the 
various  groups  and  the  relations  among  the  groups.  The 
number  of  these  questions  has  been  considerably  reduced, 
and  it  is  now  felt  even  more  strongly  than  before  that  all 
of  the  general  questions  should  be  mastered  by  every  student 
who  takes  the  course.  For  the  convenience  of  teachers  as 
well  as  of  enthusiastic  students,  a  number  of  additional  gen- 
eral questions  have  been  placed  in  the  appendix.  An  intro- 
ductory section  has  been  included,  explaining  principles  and 
details  of  laboratory  manipulation.  A  new  chapter  (Chap- 
ter VII)  has  been  added,  which  embraces  the  non-metallic 
elements,  and  is  intended  to  recall  and  to  broaden  the 
knowledge  which  the  student  was  supposed  to  possess  be- 
fore commencing  the  study  of  the  metallic  elements.  Some 
useful  tables  have  been  added  in  the  appendix,  and  a  few 
cuts  have  been  inserted  in  the  text. 

June,  1910. 


CONTENTS 

INTRODUCTORY  TO  THE  STUDENT    i 

NOTES  ON  LABORATORY  MANIPULATION 5 

1.  Precipitation-;  Crystallization 5 

2.  Pouring 6 

3.  Transferring  Precipitates  or  Crystals 6 

4.  Filtering;  Collecting  Precipitates 7 

(a)  A  Coarse-Grained  Crystal  Meal 7 

(b)  Filtering  with  Suction  ;  Witt  Filter 7 

Suction 8 

(c)  Filtering  without  Suction 9 

(d)  Filtering  Corrosive  Liquids 10 

(e)  Cloudy  Filtrates u 

(/)  To  Keep  Liquids  Hot  during  Filtration 1 1 

(g)  Cloth  Filters u 

5.  Washing  Precipitates 12 

(a)  Washing  on  the  Filter , 12 

(b)  Washing  by  Decantation 13 

6.  Evaporation      14 

7.  Dissolving  Solid  Substances      15 

8.  Crystallization 15 

9.  Drying 18 

10.  Pulverizing 19 

11.  Dry  Reactions;  Furnaces 20 

CHAPTER  I.     ALKALI  AND  ALKALINE  EARTH  METALS     ...  23 

1.  Potassium  Nitrate  from  Sodium  Nitrate  and  Potassium 

Chloride 25 

2.  Caustic  Potash  from  Wood  Ashes 28 

3.  Sodium  Carbonate  by  the  Ammonia  Process 31 

4.  Chemically  Pure  Sodium  Chloride 34 

5.  Ammonium  Bromide 37 

6.  Strontium  Hydroxide  from  Strontium  Sulphate    ....  39 

7.  Strontium  Chloride  from  Strontium  Sulphate 42 

8.  Barium  Oxide  and  Barium  Hydroxide  from  Barium  Car- 

bonate      45 

GENERAL  QUESTIONS  I   .................  47 


X  CONTENTS 

CHAPTER   II.     ELEMENTS    OF    THE    THIRD    GROUP    OF    THE 

PERIODIC  SYSTEM 51 

9.     Boric  Acid 53 

10.  Alum  from  Cryolite .    .  54 

11.  Aluminum  Sulphide 59 

GENERAL  QUESTIONS  II 61 

CHAPTER  III.     HEAVY  METALS  OF  THE  FIRST  Two  GROUPS 

OF  THE  PERIODIC  SYSTEM 63 

12.  Crystallized  Copper  Sulphate  from  Copper  Turnings  .    .  65 

13.  Cuprous  Chloride 66 

14.  Ammonium  and  Copper  Sulphate 69 

15.  Ammonio-Copper  Sulphate 71 

16.  Zinc  Oxide 74 

17.  Mercurous  Nitrate „    .    .    .    .  76 

18.  Mercuric  Nitrate 77 

19.  Mercuric  Sulphocyanate 78 

GENERAL  QUESTIONS  III 80 

CHAPTER   IV.     ELEMENTS   OF  THE   FOURTH   GROUP  OF  THE 

PERIODIC  SYSTEM 83 

20.  Stannous  Chloride 85 

21.  Stannic  Sulphide  (Mosaic  Gold) 88 

22.  Stannic  Chloride  (Anhydrous) „ 90 

23.  Lead  Nitrate 92 

24.  Lead  Dioxide 93 

25.  Red  Lead 95 

GENERAL  QUESTIONS  IV 96 

CHAPTER    V.     ELEMENTS    OF    THE    FIFTH    GROUP    OF    THE 

PERIODIC  SYSTEM 99 

26.  Ortho-Phosphoric  Acid 101 

27.  Arsenic  Acid 103 

28.  Antimony  Trichloride  from  Stibnite 106 

29.  Sodium  Sulphantimonate 108 

30.  Antimony  Pentasulphide no 

31.  Metallic  Antimony in 

32.  Basic  Bismuth  Nitrate  (Bismuth  Subnitrate) 112 

GENERAL  QUESTIONS  V 114 


CONTENTS  XI 

CHAPTER  VI.     HEAVY  METALS  OF  THE  SIXTH,  SEVENTH,  AND 

EIGHTH  GROUPS  OF  THE  PERIODIC  SYSTEM  ...  117 

33.  Potassium  Bichromate  from  Chromite 119 

34.  Potassium  Chromate  from  Potassium  Bichromate    ...  122 

35.  Chromic  Anhydride 122 

36.  Chromic  Alum 124 

37.  Chromium  Metal  by  the  Goldschmidt  Process 127 

38.  Manganese  Chloride  from  Waste  Manganese  Liquors     .  128 

39.  Potassium  Permanganate  from  Manganese  Dioxide     .    .  131 

40.  Manganese  Metal  by  the  Goldschmidt  Process      ....  134 

41.  Ferrous  Ammonium  Sulphate 135 

42.  Ferric  Ammonium  Alum 136 

GENERAL  QUESTIONS  VI 138 

CHAPTER    VII.     NON-METALLIC    ELEMENTS    OF   THE    SIXTH 

AND  SEVENTH  GROUPS  OF  THE  PERIODIC  SYSTEM,  143 

43.  Potassium  Iodide 145 

44.  Hydriodic  Acid 146 

45.  Potassium  Chlorate 149 

46.  Potassium  Bromate 153 

47.  Potassium  lodate 155 

48.  lodic  Acid;  Iodine  Pentoxide       157 

49.  Potassium  Perchlorate 159 

50.  Sodium  Thiosulphate 160 

GENERAL  QUESTIONS  VII 163 

APPENDIX 

Additional  General  Questions 165 

Atomic  Weights 176 

Periodic  Arrangement  of  the  Elements 177 

Table  of  Solubilities 178 


SYNTHETIC    INORGANIC    CHEMISTRY 


Introductory  to  the  Student 

THE  following  exercises  are  designed  to  illustrate  the 
principles  and  methods  involved  in  the  preparation  of  a 
number  of  the  most  important  chemicals.  Where  possible 
the  method  employed  resembles  that  actually  used  on  an 
industrial  scale ;  where  this  is,  however,  impossible  on  the 
limited  scale  of  the  laboratory,  mention  is  made  of  the 
fact,  with  reasons  therefor.  On  account  of  the  limitations 
connected  with  work  on  a  laboratory  scale,  it  is  of  course 
impossible  to  get  as  high  percentage  yields  as  could  be 
obtained  on  a  large  commercial  scale.  The  amounts  ob- 
tained of  each  preparation  are  to  be  weighed  and  recorded, 
but  the  chief  stress  is  to  be  laid  upon  the  excellence  of  the 
product  rather  than  upon  its  quantity. 

A  larger  number  of  preparations  is  given  than  it  is 
expected  that  the  student  can  accomplish  in  fifteen  weeks 
with  but  four  hours  per  week  in  the  laboratory.  Each 
student,  therefore,  will  be  assigned  certain  of  the  exercises 
which  he  is  expected  to  thoroughly  master,  and  which  he 
is  expected  to  perform  entirely  independently.  But  almost 
equal  in  importance  is  it  for  him  to  know  the  work  which 
the  students  about  him  in  the  laboratory  are  performing. 
To  this  end  it  is  important  that  the  directions,  and  more 
especially  the  first  sections  which  discuss  the  principles 
involved,  be  studied  for  each  exercise,  and  then  that  the 
work  of  neighboring  students  actually  at  work  upon  the 
preparations  be  observed  and  discussed  with  them  in 

I 


2  '  -  t\NTHfeTie.  INORGANIC    CHEMISTRY 

the  odd  moments  which  will  invariably  occur  when  waiting 
for  evaporations  or  filtrations  to  take  place. 

Directions  for  laboratory  work:  The  notes  for  each 
preparation  are  divided  into  three  parts: 

I.  Discussion  of  the  general  principles  involved. 

II.  Directions  for  actual  manipulation. 

III.  Study  questions. 

Part  I  is  to  be  read  and  understood  before  commencing 
work  in  the  laboratory. 

Part  II,  being  the  working  directions,  is  to  be  kept  av 
hand  while  carrying  out  the  manipulations.  These  direc- 
tions do  not  need  to  be  recorded  in  the  laboratory  note< 
book ;  but  it  is  essential,  nevertheless,  to  keep  a  laboratory 
notebook  in  which  to  enter  all  important  observations  and 
data ;  such  as,  for  example,  appearance  of  solutions  (color, 
turbidity)  ;  appearance  of  precipitates  or  crystals  (color,  size 
of  grains,  crystalline  form)  ;  results  of  all  weighings  or 
measurements ;  number  of  recrystallizations ;  results  of 
tests  for  purity  of  materials,  etc. 

Part  III  constitutes  directions  for  study  based  upon 
the  particular  preparation.  This  will  involve  :  (i)  labora- 
tory experiments  and  direct  entries  in  the  laboratory  note- 
book ;  (2)  consultation  of  reference  books,  of  which  all  that 
are  necessary  will  be  found  upon  the  shelf  in  the  laboratory ; 
(3)  original  reasoning. 

The  answers  to  the  questions  should  be  written  in  the 
laboratory  notebook  following  the  entries  for  the  exercise, 
and  this  book  should  be  submitted  at  the  same  time  as  the 
preparation  for  the  approval  of  an  instructor. 

Besides  the  specific  study  questions  for  each  prepara- 
tion there  are,  accompanying  each  group  of  exercises, 
general  questions  relating  to  the  whole  group;  and  these 
are  to  be  worked  out  by  every  student.  The  answers  to 
these  questions  are  to  be  written  on  a  certain  prescribed 


INTRODUCTORY  3 

kind  of  paper  and   handed   in  at  the  office,  neatly  bound, 
within  the  times  which  will  be  posted. 

In  preparation  work  it  is  frequently  necessary  to  wait 
for  considerable  periods  of  time  for  evaporations,  crystalliza- 
tions, etc.,  to  take  place.  This  time  may  be  utilized  for 
work  upon  the  study  questions  and  experiments,  but  even 
then  it  is  advisable  to  have  usually  more  than  a  single 
preparation  under  way.  Thus  no  time  need  be  wasted  by 
the  energetic  student  who  plans  his  work  we1!. 


NOTES    ON    LABORATORY 
MANIPULATION 

THESE  notes  are  intended  simply  to  help  the  student  in 
foreseeing  and  in  overcoming  some  of  the  difficulties  that 
arise  in  experimental  work.  They  by  no  means  make  it 
unnecessary  for  him  to  exercise  ingenuity  and  originality 
in  planning  and  carrying  out  the  details  of  laboratory  work. 
At  the  outset  these  notes  should  be  read  through  carefully ; 
then  when  in  the  later  work  references  to  specific  notes  are 
made  their  general  bearing  will  be  better  appreciated. 

i.     PRECIPITATION;    CRYSTALLIZATION 

In  the  majority  of  chemical  processes  which  are  carried 
out  in  the  wet  way,  separations  are  accomplished  by  taking 
advantage  of  differences  in  solubility.  In  case  a  certain 
product  is  extremely  insoluble  and  is  formed  almost  instan- 
taneously when  solutions  containing  the  requisite  compo- 
nents are  mixed,  the  process  is  called  precipitation  and  the 
insoluble  substance  is  called  the  precipitate.  If  the  product 
to  be  formed  is  less  insoluble,  so  that  it  separates  more 
slowly,  or  only  after  evaporating  away  a  part  of  the  solvent, 
the  process  is  called  crystallization. 

In  some  cases  the  precipitate,  or  the  crystals,  constitute 
the  desired  product ;  in  other  cases  a  product  which  it  is 
necessary  to  remove  from  the  solution  before  the  desired 
product  can  be  obtained  pure.  In  either  case  it  is  neces- 
sary to  make  as  complete  a  separation  as  possible  of  the 
solid  from  the  liquid.  This  involves  the  manipulations 
described  under  Notes  2,  3,  and  4. 

5 


O  NOTES    ON    LABORATORY    MANIPULATION 

2.      POURING 

In  pouring  a  liquid  from  a  vessel,  either  into  a  filter  or 
into  another  vessel,  care  must  be  taken  not  to  slop  the  liquid 
nor  to  allow  it  to  run  down  the  outside  of  the  vessel  poured 
from.  To  this  end  touch  a  stirring  rod  to  the  lip  of  the 


FIG.    I  FIG.    2 

dish  or  beaker  (Fig.   i)  and  allow  the  liquid  to  run  down 
the  rod. 


3.     TRANSFERRING    PRECIPITATES    OR    CRYSTALS 

If  large  crystals  have  separated  from  a  liquid  they  may 
be  picked  out,  or  the  liquid  may  be  poured  off. 

If  a  precipitate  or  a  crystalline  meal  has  formed  it  must 
be  drained  in  a  filter  funnel.  First  pour  off  the  liquid  (see 
Note  2)  —  through  the  filter  if  necessary,  so  as  to  save  any 


FILTERING  7 

floating  particles  of  the  solid  —  then  pour  the  main  part 
of  the  damp  solid  into  the  filter.  A  considerable  part  of 
the  solid  will  adhere  to  the  dish ;  most  of  this  may  be 
scraped  out  by  means  of  a  spatula,  but  the  last  of  it  is 
most  easily  rinsed  into  the  filter.  For  rinsing,  a  jet  of 
water  from  the  wash  bottle  (Fig.  2)  may  be  used  if  the  solid 
is  very  insoluble.  If  the  solid  is  soluble  in  water,  some  of 
the  saturated  solution  may  be  poured  back  into  the  dish 
from  out  of  the  filter  bottle,  and  by  means  of  this  the  last 
of  the  solid  may  be  removed  to  the  filter. 


4.     FILTERING;    COLLECTING    PRECIPITATES 

(a)  A  coarse-grained  crystal  meal  can  best  be  collected 
in  a  filter  funnel  in  which  a  perforated  porcelain  plate  is 
placed,  and  the  mother  liquor  clinging  to  the  crystals  can 

best  be  removed  with  the  aid 
of  suction  (see  next  para- 
graph). 

(fi)  Pilfering  with  Suc- 
tion;  Witt  Filter.  —  With  a 
fine-grained  crystal  meal,  or 
a  precipitate  which  is  not  of 
such  a  slimy  character  as  to 
clog  the  pores  of  the  filter 
paper,  a  Witt  filter  is  most 
advantageously  used.  For 
most  general  use  a  5-inch 
filter  funnel  should  be  fitted 
tightly  by  means  of  a  rub- 
ber stopper  into  the  neck 
of  a  500  cc.  filter  bottle 
(Fig.  3).  Place  a  i^-inch 
FIG.  3  perforated  filter  plate  in  the 


8  NOTES    ON    LABORATORY    MANIPULATION 

funnel  and  on  this  a  disk  of  filter  paper  cut  so  that  its 
edges  will  turn  up  about  3  mm.  on  the  side  of  the  funnel 
all  the  way  around.  Hold  the  disk  of  dry  paper  in  the 
right  position,  wet  it  with  a  jet  from  the  wash  bottle, 
draw  it  firmly  down  against  the  filter  plate  by  applying 
the  suction,  and  press  the  edges  firmly  against  the  side 
of  the  funnel,  so  that  no  free  channel  shall  remain.  In 
pouring  the  liquid,  direct  it  with  a  stirring  rod  (Note  2) 
onto  the  middle  of  the  filter ;  do  not  allow  it  to  run  down 
the  side  of  the  funnel,  as  this  might  turn  up  the  edge  of 
the  paper  and  allow  some  of  the  precipitate  to  pass  through. 
After  bringing  all  of  the  solid  upon  the  filter  it  may  be 
freed  from  a  large  part  of  the  adhering  liquid  by  means 
of  the  suction,  and  it  may  then  be  purified  by  washing 
with  a  suitable  liquid  (see  Note  5). 

The  Witt  filter  is  very  generally  useful  for  the  purpose 
of  separating  a  solid  product  from  a  liquid.  In  cases  that 
the  liquid  runs  slowly,  the  rate  of  filtration  can  be  increased 
by  using  a  larger  filter  plate  or  still  better  a  Biichner  fun- 
nel and  thereby  increasing  the  filtering  area.  The  student 
should,  however,  avoid  using  the  suction  indiscriminately, 
for  in  many  cases  it  is,  as  explained  in  paragraph  c,  a  positive 
disadvantage. 

Suction.  —  The  most  convenient  source  of  suction  is  the 
Richards  water  pump,  which  can  be  attached  directly  to 
the  water  tap.  If  the  water  is  supplied  at  a  pressure  of 
somewhat  over  one  atmosphere  (34  feet  of  water),  a  vacuum 
of  very  nearly  an  atmosphere  can  be  obtained.  If  the  pres- 
sure is  insufficient,  an  equally  good  vacuum  can  be  obtained 
by  means  of  the  suction  of  the  escaping  water.  To  this 
end  the  escape  pipe  must  be  prolonged  by  a  tube  sufficiently 
constricted  to  prevent  the  sections  of  the  descending  water 
column  from  breaking  and  thus  allowing  air  to  enter  from 
the  bottom. 


FILTERING  9 

To  keep  the  suction  pump  working  continuously,  how- 
ever, is  extravagant  of  water  as  well  as  being  a  nuisance 
in  the  laboratory  on  account  of  the  unnecessary  noise. 
Consequently  this  rule  is  made  and  must  be  observed: 

The  suction  pump  must  never  be  kept  in  operation  more 
than  two  minutes  at  one  time. 

If  suction  must  be  applied  for  more  than  that  length 
of  time,  the  vacuum  which  is  produced  inside  of  the  two 
minutes  may  be  maintained  in  the  suction  bottle  by  closing 
the  latter  air  tight.  For  this  purpose  the  bottle  is  to  be 
fitted  as  follows : 

Connect  a  short  piece  of  rubber  tube  with  the  side  arm 
of  the  filter  bottle.  Provide  this  tube  with  a  screw  cock 
and  connect  its  further  end  with  a  short  piece  of  glass  de- 
livery tube  tapered  a  little  at  each  end  and  rounded  in  the 
flame.  (See  Fig.  3  on  page  7.) 

The  short  glass  tube  can  be  attached  to  and  removed 
from  the  pump  at  will,  and  a  vacuum  once  produced  in  the 
bottle  can  be  preserved  by  closing  the  screw  cock.  Thus, 
for  example,  if  all  the  joints  of  the  bottle  are  tight,  a  slimy 
precipitate  may  be  left  filtering  under  suction  over  night 
or  even  longer. 

(V)  Filtering  without  Suction.  —  A  slimy  or  gelatinous 
precipitate  can  be  collected  much  better  without  suction. 
Suction  drags  the  solid  matter  so  completely  into  the  pores 
of  the  filter  that  in  most  cases  the  liquid  nearly  ceases  to 
run.  A  filter  funnel  and  filter  should  be  chosen  large 
enough  to  hold  the  entire  precipitate.  The  filter  paper 
should  be  folded  twice  and  then  opened  out  in  the  form 
of  a  cone  and  fitted  into  the  funnel  (Instructions).  The 
corners  of  the  filter  should  be  cut  off  round,  and  the 
upper  edge  of  the  filter  should  come  about  one-half  inch 
below  the  rim  of  the  funnel.  It  is  usually  best  to  fit  the 
paper  carefully  into  the  funnel,  to  wet  it  and  press  it  up 


IO  NOTES    ON    LABORATORY    MANIPULATION 

against  the  glass  all  around,  so  that  there  will  be  no  air 
channels. 

In  the  case  of  slow-running  liquids,  if  a  large  filter  is 
used,  it  may  be  filled  at  intervals  and  left  to  take  care  of 
itself  the  rest  of  the  time  while  other  work  is  being  done. 

In  case  a  considerable  weight  of  liquid  is  to  come  on  the 
point  of  the  filter,  this  may  be  reenforced  by  means  of  a  piece 
of  linen  cloth,  which  should  be  placed  under  the  middle  of 
the  filter  paper  before  it  is  folded,  and  should  then  be  folded 
in  with  it  so  as  to  strengthen  the  point. 

After  the  precipitate  is  collected  in  the  filter  and  drained, 
it  should  if  necessary  be  washed  (see  Note  5  on  page  12). 

Both  filtration  and  washing  take  place  much  more  rapidly 
if  the  liquid  is  hot.  Time  can  also  usually  be  saved  if  the 
precipitate  is  allowed  to  settle  as  completely  as  possible 
before  commencing  to  filter.  The  clear  liquid  can  then  be 
decanted  off,  or  if  necessary  poured  rapidly  through  the 
filter  before  the  latter  becomes  clogged  with  the  main  part 
of  the  precipitate. 

(d)  filtering  Corrosive  Liquids.  —  Solutions  of  very 
strong  oxidizing  agents,  concentrated  solutions  of  the  strong 
acids  and  bases,  and  concentrated  solutions  of  a  few  salts 
of  the  heavy  metals  —  notably  zinc  chloride  and  stannous 
chloride  —  attack  filter  paper  strongly.  Ordinary  paper  is 
thus  entirely  unserviceable  for  filtration,  but  a  felt  made 
of  asbestos  fibers  is  in  many  cases  very  useful.  Shredded 
asbestos,  which  has  been  purified  by  boiling  with  hydro- 
chloric acid  and  subsequent  washing,  is  suspended  in  water ; 
the  suspension  is  poured  onto  a  perforated  plate  placed 
in  a  filter  funnel ;  and  suction  is  applied  whereby  the  water 
is  removed  and  the  fibers  are  drawn  together  to  form  a 
compact  felt  over  the  filter  plate.  Enough  asbestos  should 
be  used  to  make  a  felt  i  to  3  mm.  thick,  and  care  must  be 
taken  to  see  that  it  is  of  uniform  thickness  and  that  no  free 


FILTERING  I  I 

channels  are  left  through  which  solid  matter  may  be  drawn. 
Before  it  is  ready  for  use  a  considerable  amount  of  water 
should  be  drawn  through  the  filter,  and  the  loose  fibers 
should  be  rinsed  out  of  the  filter  bottle.  Before  pouring 
the  liquid  onto  the  filter  the  suction  should  be  started 
gently,  and  the  liquid  should  be  directed  by  means  of  a 
stirring  rod  (Note  2)  onto  the  middle  of  the  filter.  If  these 
precautions  are  not  observed  the  felt  may  become  turned  up 
in  places,  so  that  the  precipitate  will  pass  through. 

(e)  Cloudy  Filtrates.  —  When    a   filtrate   at   first    comes 
through  cloudy,  it  is  usually  sufficient  to  pour  the  first  por- 
tion  through  the  filter  a  second  time.     The  pores  of  the 
filter  soon  become  partially  closed  with  the  precipitate,  so 
that  then  even  the  finest  particles  are  retained. 

With  some  very  fine-grained  precipitates,  repeatedly  pour- 
ing the  filtrate  through  the  same  filter  will  finally  give  a  clear 
filtrate. 

Special  kinds  of  filter  paper  are  made  to  retain  very  fine 
precipitates,  but  they  allow  the  liquid  to  pass  much  more 
slowly  than  ordinary  filters,  and  their  use  is  by  no  means 
essential  in  any  of  the  following  preparations. 

(f)  To   Keep   Liquids   Hot  during  Filtration.  —  When 
liquids  must  be   kept  hot  during  a  slow  filtration,  as,  for 
example,  when  cooling  would  cause  a  separation  of  crystals 
that  would  clog  the  filter,  it  sometimes  becomes  necessary 
to  surround  the  funnel  with  a  jacket  which  is  heated  with 
steam  or  boiling  water.     In  the  following  preparations  the 
use  of  such  a  device  will  not  be  necessary,  although  there 
are  several  instances  where  it  is  necessary  to  work  quickly 
to  avoid  clogging  the  filter. 

(£•)  Cloth  Filters.  —  In  preparations  made  on  a  small 
scale,  paper  filters  placed  in  ordinary  filter  funnels  are  inva- 
riably used  if  the  liquid  is  not  too  corrosive.  On  a  larger 
scale  or  in  commercial  practice,  cloth  is  much  used  for  fil- 


12  NOTES    ON    LABORATORY    MANIPULATION 

ters,  and  it  can  be  made  in  the  shape  of  bags  or  it  can  be 
stretched  over  wooden  frames.  The  cloth  or  other  filtering 
medium  (asbestos,  paper  pulp,  sand,  etc.)  has  to  be  chosen 
in  each  case  with  reference  to  the  nature  of  the  precipitate 
and  the  corrosiveness  of  the  liquid. 

Many  of  the  preparations  in  this  book,  if  carried  out  on 
a  larger  scale  than  given  in  the  directions,  would  require 
the  use  of  such  cloth  filters.  It  is  often  advantageous  to 
tack  one  piece  of  cloth  permanently  across  a  wooden  sup- 
port and  on  top  of  this  to  lay  a  second  cloth.  The  precipi- 
tate can  then  be  easily  removed  together  with  the  unfastened 
cloth. 

For  devices  for  rapid  filtration  and  filtration  in  general 
on  a  large  scale,  a  work  on  Industrial  Chemistry  should  be 
consulted. 

5.     WASHING    PRECIPITATES 

(a)  Washing  on  the  Filter.  —  To  remove  completely  the 
impurities  contained  in  the  mother  liquor  clinging  to  pre- 
cipitates or  crystals,  the  solid  is  washed.  Pure  water  is 
used  for  washing,  provided  the  solid  is  not  too  soluble  or 
is  not  decomposed  (hydrolyzed)  by  it.  Special  directions 
will  be  given  when  it  is  necessary  to  use  other  than  pure 
water. 

First,  the  solid  should  be  allowed  to  drain  as  completely 
as"  possible,  then  the  wash  liquid  should  be  applied,  prefer- 
ably from  the  jet  of  a  wash  bottle,  so  as  to  wet  the  whole 
mass  and  to  rinse  down  the  sides  of  the  filter.  If  suction 
is  used,  suck  the  solid  as  dry  as  possible,  then  stop  the 
suction  while  applying  the  washing  liquid ;  after  the  solid 
is  thoroughly  wet,  suck  out  the  liquid  and  repeat  the 
washing. 

A  little  thought  will  make  it  clear  that  the  washing  is 
much  more  effective  if  the  liquid  is  removed  as  completely 


WASHING     PRECIPITATES 


as  possible  each  time  before  applying  fresh  wash  liquid,  and 
that  a  number  of  washings  with  a  small  amount  of  liquid 
each  time  is  more  effective  than  fewer  washings  with  much 
greater  quantities  of  wash  liquid.  It  is,  of  course,  evident 
that  with  each  washing  the  liquid  should  penetrate  to  all 
parts  of  the  solid  material. 

(b)  Washing  by  Decantation.  —  In  case  a  precipitate  is 
very  insoluble  it  can  be  most  thoroughly  and  quickly  washed 
by  decantation.  This  consists  in  allowing  it  to  settle  in 
a  deep  vessel  and  then  in  pouring  (decanting)  or  siphon- 
ing off  the  clear  liquid.  Following  this  the  precipitate  is 
stirred  up  with  fresh  water  and  allowed  to  settle,  and  the 
liquid  is  again  decanted  off.  By  a  sufficient  number  of 
repetitions  of  this  process,  the  precipitate  may  be  washed 
entirely  free  from  any  soluble  impurity,  after  which  it  may 
be  thrown  on  a  filter,  drained,  and  then  dried. 

Most  precipitates,  even  after  they  have  settled  as  com- 
pletely as  possible  in  the  liquid  from 
which  they  were  thrown  down,  are  very 
bulky,  and  their  apparent  volume  is 
very  large  as  compared  with  the  ac- 
tual volume  of  the  solid  matter  itself. 
For  example,  a  precipitate  of  basic  zinc 
carbonate  (No.  16,  page  74),  after  it 
has  settled  as  completely  as  possible 
in  a  deep  jar  (Fig.  4),  may  still  occupy 
a  volume  of  400  cc.  When  this  bulky 
precipitate  is  dried,  however,  it  shriv- 
els up  into  a  few  small  lumps  whose 
total  volume  is  not  more  than  4  or  5  cc. 

If  a  precipitate,  which  is  at  first 
uniformly  suspended  in  a  liquid,  is 
allowed  to  settle  in  a  tall  jar  until  it 
occupies  but  J  of  the  original  volume  FIG.  4 


14  NOTES    ON    LABORATORY    MANIPULATION 

of  the  mixture  (Fig.  4),  any  soluble  substances  will  still 
remain  uniformly  distributed  throughout  the  whole  volume. 
If  now  the  upper  f,  consisting  of  the  clear  solution,  is 
drawn  away,  it  follows  that  practically  |-  of  the  solution, 
containing  £  of  the  soluble  impurities,  remains  with  the  pre- 
cipitate. By  stirring  up  the  solid  again  with  pure  water, 
the  soluble  impurities  become  uniformly  distributed  through 
the  larger  volume,  and  on  letting  the  precipitate  settle  and 
drawing  off  f  of  the  liquid,  as  before,  there  will  remain  with 
the  wet  precipitate  only  £  X  i  =  ^5  of  the  original  soluble 
matter.  After  the  third  decantation  the  remaining  suspen- 
sion will  contain  ^  X  sV  —  lir  °^  tne  original  impurities, 
and  so  on. 

6.     EVAPORATION 

When  it  is  necessary  to  remove  a  part  of  the  solvent 
from  a  solution,  as  when  a  dissolved  substance  is  to  be 
crystallized  out,  the  solution  is  evaporated.  In  some  cases, 
where  the  dissolved  substance  is  volatile  or  is  decomposed 
by  heat,  the  evaporation  must  take  place  at  room  temper- 
ature, but  ordinarily  in  the  following  preparations  the  liquid 
may  be  boiled.  The  boiling  down  of  a  solution  should 
always  be  carried  out  in  a  porcelain  dish  of  such  a  size 
that  at  the  outset  it  is  well  filled  with  the  liquid.  (Never 
evaporate  in  a  beaker.)  *  The  flame  should  be  applied 
directly  under  the  middle  of  the  dish  where  the  liquid  is 
deepest ;  the  part  of  the  dish  against  which  the  flame  plays 
directly  should  be  protected  with  wire  gauze.  Under  no 
circumstances  should  the  flame  be  allowed  to  play  up  over 
the  sides  of  the  dish :  first,  because,  by  heating  where  the 
dish  is  part  cooled  by  liquid  and  part  uncooled,  there  is 
great  danger  of  breaking;  second,  because  by  heating  the 
sides  the  film  of  liquid  which  creeps  up  is  evaporated 
and  the  solid  deposited  becomes  baked  hard  and  in  some 


CRYSTALLIZATION  1 5 

cases  is  decomposed.  To  prevent  the  formation  of  a  solid 
crust  around  the  edges,  which  even  at  best  will  take  place 
to  some  extent,  the  dish  should  occasionally  be  tilted  back 
and  forth  a  little,  so  that  the  crust  may  be  dissolved,  or 
loosened,  and  washed  back  into  the  middle  of  the  dish. 

While  evaporating  a  liquid  over  a  flame  it  should  be 
carefully  watched,  for  if  it  should  be  forgotten  and  evapo- 
rate to  dryness  the  dish  would  probably  break  and  the 
product  be  spoiled.  If  a  precipitate  or  crystals  separate 
from  the  liquid  and  collect  in  a  layer  at  the  bottom,  the 
dish  will  probably  break,  because  where  the  solid  prevents 
a  free  circulation  of  the  liquid  the  dish  becomes  superheated, 
and  then  when  in  any  one  place  the  liquid  does  penetrate, 
the  sudden  cooling  causes  the  porcelain  to  crack.  Usually 
when  a  solid  begins  to  separate  from  a  boiling  liquid  the 
evaporation  should-  be  stopped  and  the  liquid  left  to  crystal- 
lize. After  that  the  mother  liquor  may  be  evaporated  further 
in  a  smaller  dish. 


7.     DISSOLVING    SOLID    SUBSTANCES 

The  process  of  dissolving  solid  substances  is  hastened, 
first  by  powdering  the  substance  as  finely  as  possible,  and 
second  by  raising  the  temperature.  The  solid  and  solvent 
should  be  heated  together  in  a  porcelain  dish  (not  in  a 
beaker),  and  care  should  be  taken  to  keep  the  mixture  well 
stirred,  for  if  the  solid  should  settle  in  a  layer  on  the  bot- 
tom, that  part  of  the  dish  would  become  superheated  and 
would  be  apt  to  break  (see  last  paragraph  in  Note  6). 


8.     CRYSTALLIZATION 

(a)  A  great  number  of  pure  substances  are  capable  of 
assuming  the  crystalline  condition  when  in  the  solid  form. 


1 6  NOTES    ON    LABORATORY    MANIPULATION 

Crystals  are  bounded  by  plane  surfaces,  which  make  definite 
and  characteristic  angles  with  each  other  and  with  the  so- 
called  axes  of  the  crystals. 

The  external  form  of  a  crystal  reflects  in  some  manner 
the  shape  or  structure  of  the  individual  molecules  of  the 
substance ;  for  the  crystal  must  be  regarded  as  being  built 
up  by  the  deposition  of  layer  on  layer  of  molecules,  all  of 
which  are  placed  in  the  same  definite  spatial  relation  to 
the  neighboring  molecules. 

When  a  substance  takes  on  the  solid  form  very  rapidly 
(as  when  melted  glass  or  wax  cools)  its  molecules  do  not 
have  an  opportunity  to  arrange  themselves  in  a  regular 
order,  and  consequently  the  solid  body  is  amorphous.  The 
axes  of  the  individual  molecules  point  in  every  direction 
without  regularity,  and  consequently  the  solid  body  possesses 
no  crystalline  axes  or  planes. 

It  is  evident  from  the  above  that  the  essential  condition 
favoring  the  formation  of  perfect  crystals  is  that  the  solid 
shall  be  built  up  very  slowly.  This  is  the  only  general  rule 
which  can  be  given  in  regard  to  the  formation  of  perfect 
crystals. 

The  excellence  of  a  chemical  preparation  is  in  many 
cases  judged  largely  from  its  appearance.  The  more  uni- 
form and  perfect  the  crystals,  the  better  appearance  the 
preparation  presents. 

In  the  following  preparations  sometimes  a  pure  melted 
substance  is  allowed  to  crystallize  by  simply  cooling;  in 
such  a  case  the  cooling  should  take  place  slowly.  More 
often  crystals  are  formed  by  the  separation  of  a  dissolved 
substance  from  a  saturated  solution.  Perfect  crystals  can 
best  be  obtained  in  this  case  by  keeping  the  solution  at  a 
constant  temperature  and  allowing  it  to  evaporate  very 
slowly.  This  is  easily  accomplished  in  industrial  works 
where  large  vats  of  solution  can  be  kept  at  a  uniform  tern- 


CRYSTALLIZATION  I/ 

perature  with  steam  coils  and  allowed  to  evaporate  day  and 
night.  On  the  laboratory  scale  it  is  almost  impossible,  first 
on  account  of  variations  in  temperature,  and  next  on  account 
of  dust  which  must  fall  into  an  uncovered  dish. 

The  majority  of  substances  are  more  soluble  at  higher 
temperatures  than  at  lower.  If  a  solution  just  saturated 
at  a  high  temperature  is  allowed  to  cool  very  slowly,  it  is 
possible  for  the  solid  to  separate  so  slowly  as  to  build  up 
perfect  crystals.  This  is  an  expedient  that  can  be  adopted 
to  advantage  in  several  of  the  following  preparations.  In 
many  cases,  however,  when  a  saturated  solution  cools  it 
becomes  supersaturated,  sometimes  to  a  high  degree.  Then 
when  crystallization  is  once  induced  it  occurs  with  such 
rapidity  that  a  mass  of  minute  crystals,  instead  of  a  few 
large,  perfect  ones,  is  produced.  To  avoid  this  supersatu- 
ration  a  few  seed  crystals  (i.  e.,  very  small  crystals  of  the 
kind  desired)  may  be  placed  in  the  solution.  These  form 
nuclei  on  which  large  crystals  can  be  built  up,  and  when 
they  are  present  it  is  impossible  for  the  solution  to  remain 
supersaturated. 

In  carrying  out  the  following  preparations  the  principles 
just  stated  should  be  kept  carefully  in  mind ;  but  in  many 
instances  specific  suggestions  will  be  given  as  to  the  easi- 
est method  for  obtaining  good  crystals  of  any  particular 
substance. 

Large  crystals,  it  is  true,  present  a  pleasing  appearance, 
but  oftentimes  they  contain  a  considerable  quantity  of  the 
mother  liquor  inclosed  between  their  crystal  layers.  Hence 
if  purity  of  product  is  the  sole  requisite,  it  is  often  more 
desirable  to  obtain  a  meal  of  very  fine  crystals.  Such  a 
meal  is  obtained  by  crystallizing  rapidly  and  stirring  while 
crystallizing.  Some  substances  are  so  difficult  to  obtain  in 
large  crystals  that  it  is  more  satisfactory  to  try  only  to 
obtain  a  uniform  crystal  meal. 


1 8  NOTES    ON    LABORATORY    MANIPULATION 

(£)  Purification  by  Recrystallization.  —  When  a  given 
substance  crystallizes  from  a  solution,  it  most  generally 
separates  in  a  pure  condition  irrespective  of  any  other  dis- 
solved substances  the  solution  may  contain.  Thus  a  sub- 
stance can  be  obtained  in  an  approximate  state  of  purity 
by  a  single  crystallization.  Portions  of  the  mother  liquor 
(containing  dissolved  impurities)  are,  however,  usually  en- 
trapped between  the  layers  of  the  single  crystals,  not  to 
mention  the  liquid  which  adheres  to  the  crystal  surfaces. 
By  dissolving  the  crystals,  the  small  amount  of  impurity 
likewise  passes  into  the  solution,  but  only  a  small  fraction 
of  this  impurity  is  later  entrapped  by  the  crystals  when  they 
separate  from  this  new  mother  liquor.  By  several  recrys- 
tallizations,  then,  a  substance  can  ordinarily  be  obtained 
m  a  very  high  state  of  purity. 


9.     DRYING 

(a)  A  preparation  that  is  not  affected  by  the  atmosphere 
can  be  dried  by  being  spread  in  a  thin  layer  on  a  plate 
of  glass,  on  filter  paper,  or,  best  of  all,  on  an  unglazed 
porous  porcelain  plate.1  The  substance  may  with  advan- 
tage be  turned  over  occasionally  with  a  spatula ;  and  if  it 
is  not  decomposed  by  heat  it  can  be  dried  more  rapidly  in 
a  warm  place,  such  as  over  the  steam  table.  A  product  con- 
taining water  of  crystallization  should  never  be  dried  at  an 
elevated  temperature.  During  the  drying  the  preparation 
must,  of  course,  be  carefully  protected  from  dust. 

(V)  Substances  which  decompose  on  standing  exposed  to  the 
air  may  be  quickly  dried  if  they  are  first  rinsed  with  alcohol, 
or  with  alcohol  and  then  ether.  Rinsing  with  alcohol  re- 
moves nearly  all  of  the  adhering  water,  and  a  further  rinsing 

1  Dishes  which  are  imperfect,  and  on  that  account  have  not  been 
glazed,  can  be  obtained  very  cheaply  from  the  factories. 


PULVERIZING  19 

with  ether  removes  the  alcohol.  Alcohol  evaporates  more 
rapidly  than  water,  but  ether  evaporates  so  rapidly  that  a 
preparation  wet  with  it  may  be  dried  by  a  very  few  minutes' 
exposure  to  the  air. 

Alcohol  and  ether  are  both  expensive  and  should  be 
used  sparingly.  They  can  be  used  most  effectively  as  fol- 
lows :  After  all  the  water  possible  has  been  drained  from 
the  preparation,  transfer  the  latter  to  an  evaporating  dish 
and  pour  over  it  enough  alcohol  to  thoroughly  moisten  it; 
stir  it  with  a  spatula  until  the  alcohol  has  penetrated  to 
every  space  between  the  crystal  grains,  then  pour  off,  or 
drain  off,  the  alcohol  and  treat  the  preparation  in  like  man- 
ner with  another  portion  of  fresh  alcohol.  After  that  wash 
it  once  or  twice  with  ether  in  exactly  the  same  manner.  If 
it  is  necessary  to  wash  the  preparation  on  the  filter,  drain 
off  the  water  as  thoroughly  as  possible,  stop  the  suction, 
add  just  enough  alcohol  to  moisten  the  whole  mass,  and 
after  letting  it  stand  a  few  moments  drain  off  the  liquid 
completely.  Apply  a  second  portion  of  alcohol  and  portions 
of  ether  in  the  same  manner. 


10.     PULVERIZING 

In  chemical  reactions  in  which  solid  substances  are 
involved  the  action  is  limited  to  the  surface  of  the  solid, 
and  for  this  reason  it  is  evident  that  it  must  be  much 
slower  than  reactions  which  take  place  between  dissolved 
substances ;  it  is  also  evident  that  the  more  finely  powdered 
a  solid  substance,  the  greater  is  its  surface,  and  therefore 
the  more  rapidly  it  will  react. 

Most  solid  raw  materials  for  the  following  preparations 
are  supplied  in  the  powdered  form,  but  they  are  rarely 
powdered  finely  enough ;  they  should  in  general  be  further 
pulverized  until  they  no  longer  feel  gritty  beneath  the  pestle 
or  between  the  fingers. 


2O  NOTES    ON    LABORATORY    MANIPULATION 

For  grinding  any  quantity  of  a  substance  a  large  porce- 
lain mortar  (say  8  inches  in  diameter)  with  a  heavy  pestle 
is  preferable  to  the  small  mortars  usually  supplied  in  the 
desks.  One  or  more  such  mortars  is  placed  in  the 
laboratory  for  general  use. 

If  a  hard  substance  can  be  obtained  only  in  large  pieces, 
it  should  first  be  broken  with  a  hammer,  then  crushed  into 
small  particles  in  an  iron  or  steel  mortar,  after  which  it 
is  to  be  ground  in  the  porcelain  mortar.  In  the  final 
grinding  it  is  often  advisable  to  sift  the  fairly  fine  from 
the  coarser  particles,  then  to  finish  grinding  the  former  by 
itself  and  to  crush  and  grind  the  coarser  particles  apart. 


DRY    REACTIONS:    FURNACES 

Dry  solid  substances  do  not  react  appreciably  with  each 
other  at  ordinary  temperature.  Reactions  are  made  possible 
in  two  ways :  first,  the  wet  way,  in  which  the  substances 
are  dissolved  and  thus  brought  into  most  intimate  contact. 
In  many  cases  solution  also  produces  ionization,  which,  as 
is  known,  greatly  increases  chemical  activity. 

Reactions  in  the  dry  way  are  rendered  possible  by  heat. 
Heat  alone  increases  the  rapidity  of  a  chemical  reaction,  it 
being  a  general  law  that  the  speed  is  increased  from  two 
to  three  times  for  every  increase  of  10°  C.  in  temperature. 
In  cases  in  which  one  or  more  of  the  reacting  substances 
are  melted  by  the  heat,  the  same  sort  of  intimate  contact 
is  brought  about  as  in  the  case  of  solutions.  Fusion  is 
likewise  a  means  of  producing  electrolytic  dissociation,  and 
on  this  account  also  it  increases  chemical  activity. 

In  some  of  the  furnace  reactions  in  which  none  of  the 
substances  are  melted,  as,  for  example,  in  the  reduction  of 
strontium  sulphate  to  strontium  sulphide  by  means  of  char- 
coal (see  Preparation  No.  6),  the  process  probably  takes 


FURNACES 


21 


place  in  virtue  of  a  certain  amount  of  gas  which  is  con- 
tinuously regenerated.  A  little  of  the  hot  charcoal  is 
oxidized  to  carbon  monoxide,  which  then  reduces  some  of 
the  strontium  sulphate,  it  being  itself  changed  to  carbon 
dioxide  thereby ;  the  latter  gas  comes  in  contact  with 
incandescent  charcoal,  and  carbon  monoxide  is  again 
produced. 

Reactions  in  the  dry  way  are  usually  carried  out  in 
crucibles  —  of  iron,  clay,  or  graphite,  according  as  to  which 
is  least  attacked  by  the  reagents.  For  rather  moderate  tem- 
peratures the  crucible 
may  be  heated  over  a 
flame,  but  in  most  cases 
the  requisite  temperature 
can  best  be  obtained  in 
a  furnace. 

The  form  of  furnace 
most  to  be  recommended 
for  this  work  is  repre- 
sented in  Figure  5.  It 
consists  of  a  cylinder  of  FIG.  5 

fire    clay,    7    inches    high 

and  6^  inches  in  external  diameter,  which  is  surrounded 
by  a  sheet  iron  casing.  It  is  heated,  as  shown,  by  a  gas- 
wind  flame,  introduced  through  an  opening  in  the  lower 
part  of  one  side.  If  a  suitable  air  blast  is  not  available,  a 
gasoline  blowpipe  (such  as  is  commonly  used  by  plumbers) 
is  almost  equally  serviceable. 

When  such  a  furnace  as  that  described  is  heated  as  hot 
as  possible  with  a  well-regulated  mixture  of  gas  and  air,  a 
temperature  of  about  1,350°  can  be  obtained.  For  carrying 
out  ordinary  chemical  preparation  work  an  accurate  enough 
measure  of  the  temperature  is  given  by  the  color  of  the 
glowing  interior  of  the  furnace,  and  the  approximate  cen- 


22  NOTES    ON    LABORATORY    MANIPULATION 

tigrade  values    corresponding    to    different    colors    are    as 
follows : 

Incipient  red  heat      .         .         .         .         '55°° 

Dull  red  heat 650° 

Red  heat 800° 

Bright  red  heat 1,000° 

Yellow  heat        ......    1,200° 

White  heat i>35o° 


CHAPTER  I 

ALKALI  AND  ALKALINE  EARTH  METALS 

These  metals  constitute  the  left  hand  or  A  families  in 
the  first  two  groups  of  the  periodic  classification  of  the 
elements,  as  shown  in  the  table  which  appears  in  the  ap- 
pendix, and  which  is  also  placed  for  convenience  inside 
the  front  cover  of  the  book. 

The  metals  of  these  two  families  are  studied  together 
because  they  are  the  extremely  active  base-forming  elements. 
On  account  of  their  great  activity  they  are  never  found 
uncombined  in  nature,  and  it  is  only  by  the  aid  of  the 
most  powerful  reducing  agencies  (for  example,  by  electroly- 
sis of  their  molten  salts)  that  the  metals  themselves  are 
extracted  from  their  compounds. 

The  alkali  metals  are  monovalent.  Their  hydroxides, 
MOH,  are  extremely  soluble  and  are  highly  dissociated  as 
bases ;  on  account  of  the  corrosive  properties  of  the  latter 
they  are  known  as  the  caustic  alkalies  —  hence  the  designa- 
tion, alkali  metals.  The  compounds  of  the  alkali  metals 
are,  with  a  very  few  exceptions,  soluble  in  water,  and  they 
are  all  strong  electrolytes. 

The  radical  ammonium,  NH4,  is  classed  with  the  alkali 
metals  on  account  of  its  ability  to  form  the  same  kinds  of 
compounds. 

The  alkaline  earth  metals  are  divalent ;  their  hydroxides, 
M(OH)2,  are  less  soluble  than  those  of  the  alkali  metals, 
but  are  nevertheless  very  strongly  basic.  The  compounds 
of  these  metals  are  not  so  generally  soluble  as  those 
of  the  alkali  metals,  and  in  particular  the  carbonates  and 
sulphates  are  mostly  insoluble. 

23 


I.     POTASSIUM  NITRATE  FROM  SODIUM  NITRATE 
AND  POTASSIUM  CHLORIDE 

The  most  important  source  of  nitrates  is  Chili  saltpeter, 
sodium  nitrate.  This  is  not  suited  for  use  in  explosives  ^n 
account  of  its  property  of  attracting  moisture  and  rendering 
the  explosive  preparation  damp.  Potassium  nitrate  is  not 
open  to  this  objection,  and  hence  large  quantities  of  it  are 
prepared,  using  sodium  nitrate  as  a  source  of  the  nitrate 
radical. 

When  two  ionizable  salts  are  dissolved  in  water  the 
resulting  solution  will  contain,  besides  the  undissociated 
molecules  and  the  ions  of  these  two  salts,  also  the  undis- 
sociated molecules  of  the  two  new  salts  which  form  by  the 
interaction  of  the  ions  present.  Which  of  these  four  salts 
will  crystallize  first  out  of  solution  depends  upon  their 
relative  solubility.  Thus  if  sodium  nitrate  and  potassium 
chloride  are  dissolved  together  in  water  the  resulting  solu- 
tion will  contain  Na+,  K+,  NO8~,  and  Cl~-ions,  together 
with  undissociated  molecules  of  NaNO8,  NaCl,  KNO3,  KC1. 
The  solubility  of  some  salts  varies  very  much  with  the 
temperature,  while  that  of  other  salts  varies  very  little. 
This  is  seen  from  the  following  table  and  diagram,  and 
practical  use  is  made  of  these  facts  by  crystallizing  succes- 
sively the  two  different  salts  at  different  temperatures. 

GRAMS  OF  SALT  SOLUBLE  IN  100  GRAMS  OF  WATER 

At  iob  At  100° 

KNO.3                      21  246 

NaCl                        36  40 

KC1                         31  56 

NaNO3                    8 1  1 80 
25 

7U 


26 


ALKALI  AND  ALKALINE  EARTH  METALS 


KN03 


50° 
Temperature/ 


100° 


Procedure,  —  Dissolve  100  grams  of  sodium  nitrate  and 
88  grams  of  potassium  chloride  in  200  cc.  of  water  and 
evaporate  in  a  porcelain  dish  to  half  that  volume.  Without 
letting  the  liquor  cool,  separate  it  from  the  crystals  which 
have  formed  during  the  evaporation.  This  is  best  accom- 
plished with  the -aid  of  suction  (see  Note  4  (b)  on  page  7, 
Witt  filter).  The  liquid  is  poured  through  the  filter  and 
then  the  crystals  are  thrown  upon  the  plate  and  pressed 
with  a  spatula,  while  applying  gentle  suction  in  order  to 
remove  as  much  as  possible  of  the  liquid  clinging  to  them. 
Pour  the  filtrate  into  a  beaker  and  set  it  aside  to  cool ;  then 
examine  the  crystals  left  on  the  filter  and  convince  yourself 


POTASSIUM    NITRATE  2 7 

that  they  consist  in  the  main  of  sodium  chloride.  (Examine 
with  a  microscope.  The  crystals  should  be  cubical.  Com- 
pare the  taste  with  that  of  pure  sodium  chloride  and  that 
of  pure  potassium  nitrate.)  By  means  of  running  tap  water 
cool  the  filtrate  to  about  10°,  and  then  separate  the  crys- 
tals of  potassium  nitrate  from  the  liquid  in  the  same  man- 
ner as  above  (see  Note  3  on  page  7,  last  sentence).  The 
filtrate  from  these  crystals  is  saturated  with  both  sodium 
chloride  and  potassium  nitrate,  and  the  larger  part  of  the 
latter  should  be  saved.  Evaporate  the  solution  in  a  smaller 
dish  until  a  considerable  quantity  of  sodium  chloride  crys- 
tallizes from  the  boiling  liquid.  Filter  hot,  as  above,  and 
crystallize  potassium  nitrate  from  the  filtrate  by  cooling. 
Unite  this  crop  of  potassium  nitrate  crystals  with  the  first. 
Test  a  very  small  portion  of  them  for  sodium  chloride  by 
dissolving  about  o.i  gram  in  2  cc.  of  water  and  adding  a 
drop  of  silver  nitrate  solution.  They  are  not  pure  and  must 
be  purified.  Weigh  roughly  the  crystals  while  still  moist, 
and  dissolve  them  in  from  half  to  three-quarters  of  their 
weight  of  hot  water.  Cool  and  separate  the  crystals  from 
the  mother  liquor.  The  latter  should  now  contain  nearly 
all  of  the  sodium  chloride  which  was  mixed  with  the  first 
crop  of  crystals.  Test  as  above  to  see  if  this  crop  is  free 
from  sodium  chloride.  If  not,  repeat  the  recrystallization 
as  many  times  as  is  .  necessary  to  get  a  perfectly  pure 
product.  A  little  of  this  should  when  dissolved  give  no 
turbidity  with  silver  nitrate  solution,  and  when  held  in  the 
flame  on  a  platinum  wire  should  color  it  the  violet  color  of 
potassium,  with  none  of  the  yellow  sodium  color.  Spread 
the  preparation  on  an  unglazed  porcelain  plate  and  allow  it 
to  dry  by  standing  exposed  to  the  air;  then  put  up  the 
salt  in  a  test  tube  or  a  small  bottle,  and  label  it  neatly. 

Questions 

i.     Define  metathesis, 


28      ALKALI  AND  ALKALINE  EARTH  METALS 

2.  When  a  metathetical  reaction  is  carried  out  in  the 
wet  way,  why  is   the   solubjlity  of  the  substances   involved 
of  importance?     Explain   why,   according  to  this   point  of 
view,  the  reactions  AgNO3  -\-  KC1  =  AgCl  -f  KNO3  and 
BaCl2  -f"  Na2SO4  =  BaSO4  -|-  2NaCl  are  much  more  com- 
plete than  the  reaction  NaNO3  +  KC1  =  KNO3  -f  NaCl. 

3.  Explain  why  fewer   operations    should   be   required 
to  prepare  potassium  nitrate  from   potassium  sulphate  and 
barium  nitrate  than  by  the  foregoing  procedure. 

2.     CAUSTIC    POTASH    FROM    WOOD    ASHES 

Of  the  mineral  constituents  of  plants,  potassium  salts 
form  an  important  part,  and,  so  far  as  these  are  salts  of 
organic  acids,  they  are  converted  into  potassium  carbonate 
when  the  plant  is  burned.  On  an  average,  wood  ashes 
contain  about  10  per  cent,  of  potassium  carbonate,  and 
before  the  advent  of  the  Leblanc  Soda  Process  this  was 
almost  the  sole  supply  of  alkali.  Even  after  this  process 
came  into  general  use,  by  which  sodium  carbonate  could 
be  obtained  from  common  salt,  wood  ashes  remained  for 
some  time  the  important  source  of  potassium  carbonate. 
In  recent  years,  however,  the  greater  part  of  the  production 
of  potassium  carbonate  has  been  derived  by  the  Leblanc 
Process  from  potassium  salts  found  in  deposits  in  the  earth, 
principally  at  Stassfurt,  Germany. 

Potassium  carbonate  being  the  principal  soluble  con- 
stituent of  wood  ashes,  it  is  extracted  with  water;  but  the 
extract  so  obtained  contains,  as  well,  the  other  soluble 
mineral  constituents,  and  also  a  considerable  amount  of 
tarry  coloring  matter  which  was  not  destroyed  in  the  com- 
bustion of  the  wood.  This  tarry  matter  is  destroyed  by 
calcination  of  the  residue  obtained  on  evaporating  the 
aqueous  extract,  and  the  calcined  mass  is  what  is  known 


POTASH    FROM    WOOD    ASHES  29 

as  crude  potash.  A  better  grade  of  commercial  potash  can 
be  obtained  by  dissolving  this  mass  in  water,  filtering,  and 
evaporating  the  solution. 

In  order  to  obtain  potassium  hydroxide  or  caustic  potash 
from  potassium  carbonate  the  solution  of  the  latter  is 
treated  with  milk  of  lime  (calcium  hydroxide).  With  this 
it  interacts,  yielding  insoluble  calcium  carbonate  and  soluble 
potassium  hydroxide, 

Ca(OH)2  +  K2C03  =  CaC03  +  2KOH. 

Procedure.  —  Tie  a  piece  of  cloth  over  the  mouth  of 
a  thistle  tube  and  insert  it  beneath  the  surface  of  a  layer 
of  sand,  one-half  inch  deep,  in  the  bottom  of  a  tall  2-liter 
bottle.  Mix  i  kilogram  of  wood  ashes  with  800  cc.  of 
hot  water  in  a  pail,  and  transfer  the  moist  mass  to  the 
bottle.  Pour  200  cc.  more  of  hot  water  over  the  surface 
of  the  ashes.  Connect  the  thistle  tube  with  a  siphon  and 
draw  off  as  much  liquid  as  possible,  perhaps  200  cc.,  using 
suction  and  drawing  the  liquid  into  a  suction  bottle  if  it 
does  not  otherwise  run  rapidly  enough.  Commence  evap- 
orating this  liquid  in  an  8-inch  evaporating  dish ;  then  pour 
300  cc.  of  hot  water  on  top  of  the  ashes  and  stir  around 
the  surface  layer.  Again  draw  off  about  300  cc.  of  liquid 
and  add  it  to  that  in  the  evaporating  dish,  and  repeat  the 
operation  until  the  liquid  drawn  off  is  nearly  colorless. 
Not  more  than  2  to  2\  liters  need  be  drawn  off  in  all. 
When  the  liquid  is  all  evaporated  remove  the  dry  residue 
to  an  iron  dish  and  heat  it  strongly  with  a  Bunsen  burner 
to  destroy  the  tarry  matter.  More  than  a  moderate  red 
heat  should  not  be  exceeded,  as  it  is  not  desirable  to  fuse 
the  salt,  but  the  heating  should  be  continued  until  the  ash 
is  white  or  at  most  contains  only  black  specks  of  completely 
charred  carbon.  When  cooled,  weigh  the  material ;  assum- 
ing it  to  consist  wholly  of  potassium  carbonate,  calculate 


3<3  ALKALI    AND    ALKALINE    EARTH    METALS 

the  amount  of  quicklime  (calcium  oxide)  necessary  to 
react  with  it  according  to  the  equation, 

CaO  +  H20  +  K2C03  =  2KOH  +  CaCO3. 

Slake  20  per  cent,  more  than  that  amount  of  lime  by 
covering  it  in  a  porcelain  dish  with  water  and  quickly 
pouring  off  the  excess  of  water.  If  the  lime  is  of  suitable 
quality  it  will  soon  grow  hot  and  crumble  to  a  powder, 
Ca(OH)2.  Take  sufficient  water  to  make  ten  times  the 
weight  of  the  crude  potash.  Stir  up  the  slaked  lime  with 
half  of  it,  thus  making  milk  of  lime.  Dissolve  the  potash 
in  the  other  half,  bring  it  to  boiling,  and  add  the  milk  of 
lime  with  stirring.  Let  the  mixture  boil  for  15  minutes 
and  then  filter,  using  a  suction  bottle  (see  Note  4  (£),  Witt 
filter).  Measure  the  volume  of  the  solution  of  caustic 
potash  obtained  and  preserve  it  in  a  stoppered  bottle. 

Test  the  strength  of  the  solution.  Measure  15  cc.  with 
a  pipette  into  a  beaker  and  add  a  drop  of  litmus  solution. 
Run  into  this  from  a  burette  a  solution  of  normal  hydro- 
chloric acid  (36.5  grams  per  liter),  drop  by  drop,  until  the 
color  just  changes  from  blue  to  red.  If  the  right  point  is 
overstepped  begin  again  with  a  fresh  sample  of  the  solu- 
tion. From  the  amount  of  acid  taken  to  neutralize  the 
sample,  calculate  the  amount  of  KOH  obtained  from  the 
wood  ashes.  Preserve  the  solution  in  a  bottle  labeled  with 
the  number  of  cubic  centimeters  of  the  solution,  with  its 
strength  in  mols  per  liter  of  potassium  hydroxide,  and 
with  the  actual  amount  in  grams  of  the  potassium  hydroxide. 

Questions 

i.  The  calcium  hydroxide  used  in  causticizing  the 
potash  is  a  slightly  soluble  solid  suspended  in  water,  its 
solubility  being  1.7  grams  'per  liter.  Explain  how,  in  spite 
of  its  limited  solubility,  the  required  amount  can  enter 
into  reaction, 


SODA    BY    THE    AMMONIA    PROCESS  3  I 

2.  Explain  why  the  caustic    potash   solution   obtained 
contains  practically  no  calcium  ions,  even  in  case  an  excess 
of  calcium  hydroxide  may  have  been  used  for  causticizing. 

3.  Analyses  of  the  crude  potash  obtained  from  various 
grades  of  wood  ashes  have  given  results  which  fall  within 
the  limits  given  in  the  table: 

K2CO3        .         ,  .  38-78  per  cent. 

Na2CO3      .         .  .  0-12  per  cent. 

K2SO4         .         .  .  13.5-40.5  per  cent. 

KC1            .         .  .  0.9-10.0  per  cent. 

Insoluble  matter  .  0.1-9.2  per  cent. 

What  substances  other  than  KOH  would  you  expect  then  to 
be  present  in  the  caustic  potash  solution  which  you  have 
prepared  ?  Look  up  the  solubility  of  calcium  sulphate  and 
calcium  hydroxide,  and  decide  whether,  in  the  presence  of 
a  large  amount  of  potassium  hydroxide,  milk  of  lime  would 
react  with  a  small  amount  of  potassium  sulphate  according  to 
the  equation : 

K2S04  +  Ca(OH)2  =  CaS04  +  2KOH. 

4.  If   a  solid   substance  crystallizes   from   the  caustic 
potash  solution    after   it   has  stood,  decide  what   it   is   by 
consulting  the  table  in  Question  3  and  a  solubility  table. 


3.     SODIUM    CARBONATE    BY    THE   AMMONIA 
PROCESS 

The  principle  employed  in  the  manufacture  of  sodium 
carbonate  from  sodium  chloride  by  the  Solvay  Process  is 
exceedingly  simple.  It  depends  primarily  upon  the  fact 
that  sodium  acid  carbonate  is  but  sparingly  soluble  in 
water ;  this  compound  is  produced  by  the  interaction  of  so- 


32      ALKALI  AND  ALKALINE  EARTH  METALS 

dium  chloride,  in  a  saturated  salt  solution,  with  ammonium 
acid  carbonate, 

(i)  Nad  +  NH4HCO3  ^  NaHCO3  +  NH4C1. 

Since  ammonia  and  salts  of  ammonium  are  very  much  more 
expensive  than  sodium  carbonate,  it  is  evident  that  the 
process  can  be  of  no  commercial  value  unless  ammonia 
can  be  recovered  and  used  again.  This  is  accomplished 
in  fact  by  treating  the  mother  liquor,  after  separation  from 
the  sodium  bicarbonate,  with  calcium  hydroxide, 

(2)  Ca(OH)2  -f  2NH4C1  =  CaCl2  +  2NH3  +  2H2O. 

In  practice  the  process  is  usually  carried  out  as  follows : 

A  nearly  saturated  salt  solution  is  purified  of  iron,  mag- 
nesia, lime,  etc.,  which  would  otherwise  get  into  the  final 
product;  there  is  then  passed  into  it  ammonia  gas  until  it 
has  absorbed  60  to  70  grams  of  NH3  per  liter,  whereupon 
carbon  dioxide  is  passed  in  until  it  has  reacted  with  the 
ammonia  to  form  ammonium  bicarbonate, 

NH4OH  -f-  C02  =  NH4HC03, 

which  in  turn  reacts  with  the  salt  according  to  (i).  Exten- 
sive precautions  necessarily  have  to  be  observed  that  prac- 
tically no  ammonia  escape  during  the  process,  so  that  the 
entire  amount  may  be  used  over  and  over  again.  It  is 
also  essential  that  as  little  carbon  dioxide  shall  be  wasted 
as  possible.  Thus  the  carbon  dioxide  is  utilized  which  is 
produced  in  converting  limestone  into  quicklime, 

CaCO3  ->  CaO  +  CO2, 

and  in  converting  sodium  bicarbonate  into  sodium  carbonate, 
2NaHC03  -»  Na2C03  +  H2O  +  CO2. 

Procedure.  —  To  50  cc.  of  concentrated  ammonium  hy- 
droxide (sp.  gr.  0.90)  add  150  cc.  of  water.  Place  in  a  flask 


SODA    BY    THE    AMMONIA    PROCESS  33 

and  add  60  grams  of  table  salt  free  from  lumps.  Shake 
until  the  salt  is  nearly  or  quite  dissolved  and  filter  the 
solution  if  it  is  not  perfectly  clear.  Pass  a  delivery  tube 
through  one  hole  of  a  double-boreal,  tightly  fitting  stopper 
placed  in  a  300  cc.  flask.  Provide  a  plug  for  the  other  hole. 
Let  the  tube  dip  into  the  solution  which  is  placed  in  the 
flask,  and  pass  in  carbon  dioxide  gas  from  a  Kipp  generator 
until  all  the  air  has  been  displaced  from  the  flask ;  then 
close  the  flask  and  allow  the  gas  to  pass  in  as  fast  as  it 
will  be  absorbed.  Occasionally,  as  the  action  seems  to 
slacken,  loosen  the  plug  for  a  moment.  Shake  the  flask 
frequently.  It  will  take  several  hours  for  the  solution  to 
absorb  sufficient  carbon  dioxide,  and  it  may  be  left  over 
night  connected  with  the  generator.  When  no  more  gas 
can  be  absorbed  pour  the  mixture  from  the  flask  upon 
a  Witt  filter  (see  Note  4  (£)).  Apply  suction  to  remove  the 
liquid  from  the  sodium  bicarbonate.  Wash  the  product 
three  times  with  15  cc.  of  ice  water  (see  Note  5  (a),  on 
page  12),  sucking  it  free  from  liquid  each  time.  Spread 
the  preparation  on  an  unglazed  plate  and  leave  it  until  it 
ceases  to  smell  of  ammonia.  Test  the  preparation  for  chlo- 
rides (see  Question  i),  of  which  it  should  not  contain  more 
than  a  trace. 

Questions 

1.  In    testing   the    preparation    for   chlorides,   the    test 
solution    must   be    acidified    with    nitric   acid    before    silver 
nitrate   is   added.     What  other  silver   salt  would  otherwise 
be    precipitated  ?     How    does   the    presence  of    nitric    acid 
prevent  this  ? 

2.  How    is    sodium  carbonate    prepared    from    sodium 
bicarbonate  ? 

3.  What  is  an  acid  salt  ?     How  does  a  solution  of  an 
acid  salt  such  as  KHSO4  behave  towards  litmus  ?     Test  the 


34      ALKALI  AND  ALKALINE  EARTH  METALS 

behavior  of  solutions  of  NaHCO3  and  of  Na2CO8  towards 
litmus.     Explain  the  cause  of  this  behavior. 

4.  Why  cannot  potassium  carbonate  be  prepared  from 
potassium  chloride  by  the  ammonia  process?  (Look  up  the 
solubility  of  potassium  bicarbonate.)  What  process  may 
be  used  to  obtain  potassium  carbonate  from  this  source  ? 


4.     CHEMICALLY    PURE    SODIUM    CHLORIDE 
FROM    ROCK    SALT 

Common  rock  salt  may  contain  other  than  sodium  chlo- 
ride up  to  10  per  cent,  of  matter,  which  consists  in  the 
main  of  the  sulphates  and  chlorides  of  potassium,  calcium, 
and  magnesium,  not  to  mention  a  considerable  amount  of 
dirt  and  insoluble  matter.  For  most  commercial  purposes 
these  impurities  are  not  harmful.  By  careful  crystallization 
of  the  salt  from  solution,  a  product  sufficiently  free  from 
these  impurities  can  be  obtained  to  be  used  as  table  salt. 
To  obtain  chemically  pure  sodium  chloride,  however,  more 
elaborate  precautions  must  be  taken.  A  satisfactory  method 
depends  upon  the  insolubility  of  sodium  chloride  in  a  con- 
centrated solution  of  hydrochloric  acid.  A  nearly  saturated 
solution  of  the  rock  salt  is  prepared,  and,  without  removing 
the  dirt  and  insoluble  matter,  enough  pure  sodium  carbonate 
is  added  to  precipitate  the  calcium  and  magnesium  in  the 
solution  as  carbonates.  Into  the  clear  filtrate  is  then  passed 
gaseous  hydrochloric  acid  until  the  greater  part  of  the 
sodium  chloride  is  precipitated,  while  the  small  amounts  of 
sulphates  and  of  potassium  salts  remain  in  the  solution. 
The  precipitate  is  drained  and  washed  with  a  solution  of 
hydrochloric  acid  until  the  liquid  clinging  to  the  crystals 
is  entirely  free  from  sulphates. 

Procedure.  —  Dissolve  25  grams  of  rock  salt  in  75  cc. 
of  water,  hastening  the  action  with  gentle  heating.  To  the 


PURE    SODIUM    CHLORIDE  35 

solution  add  about  i  gram  of  sodium  carbonate  dissolved 
in  a  few  cubic  centimeters  of  water.  Stir,  let  settle,  and 
add  a  few  drops  more  of  sodium  carbonate  solution,  and  if 
no  fresh  precipitate  is  produced  in  the  clear  part  of  the 


FIG.  6 


solution  no  more  need  be  added ;  otherwise  enough  more 
must  be  added  to  produce  this  result.  Filter  the  solution,  hot, 
through  an  ordinary  filter  (Note  4  (V)).  Prepare  pure  gase- 
ous hydrochloric  acid  as  follows  :  Place  50  grams  of  dry  rock 


36      ALKALI  AND  ALKALINE  EARTH  METALS 

salt  in  a  round-bottom  liter  flask  provided  with  a  rubber 
stopper  with  two  holes,  through  which  pass  a  thistle  tube 
reaching  to  the  bottom  of  the  flask,  and  an  exit  tube  just 
coming  through  the  stopper.  Provide  a  wash  bottle  for 
the  gas  as  follows:  A  300  cc.  bottle  is  provided  with  a 
stopper  with  three  holes.  Through  one  passes  a  glass  tube 
reaching  to  the  bottom  for  entrance  of  the  gas;  through 
another  a  thistle  tube  reaching  to  the  bottom  for  use  as 
a  safety  tube;  and  through  the  third  an  exit  tube  just 
coming  through  the  stopper.  Pour  into  the  bottle  enough 
concentrated  hydrochloric  acid  (sp.  gr.  1.2)  to  come  up 
three-quarters  of  an  inch  on  the  two  lower  tubes.  Connect 
the  exit  tube  with  a  2-inch  filter  funnel  for  delivering  the 
gas  into  the  sodium  chloride  solution.  Use  entirely  glass 
tubing,  and,  where  connections  must  be  made  with  rubber, 
bring  the  ends  of  the  glass  tubes  close  together.  Before 
commencing  to  use  this  apparatus  it  must  be  approved 
by  an  instructor.  Pour  the  sodium  chloride  solution  into  a 
beaker  of  3  inches  diameter,  and  insert  the  mouth  of  the 
funnel  below  the  surface  of  the  liquid.  Pour  gradually 
95  cc.  of  concentrated  sulphuric  acid  into  the  generating 
flask,  and  when  the  first  action  has  ceased  warm  very 
gently.  There  will  be  a  great  deal  of  frothing,  but  the 
froth  should  at  no  time  be  allowed  to  get  over  into  the 
wash  bottle.1  When  hydrochloric  acid  gas  ceases  to  be  gen- 
erated, separate  the  precipitated  sodium  chloride  from  the 
mother  liquor  by  pouring  it  upon  a  Witt  filter.  Suck 
the  moisture  from  the  crystals.  Test  the  filtrate  for  sul- 
phate by  adding  a  little  barium  chloride  solution  to  a 
small  sample  of  it  diluted  with  water.  A  strong  test  will 
probably  be  obtained.  Now  wash  the  crystals  with  succes- 

1  The  melted  sodium  acid  sulphate  left  in  the  generating  flask  is 
very  hot  and  must  not  be  poured  into  the  sink.  It  may  be  poured  into 
some  dry  receptacle  specially  provided,  or  it  may  be  allowed  to  cool 
slowly  and  solidify  in  the  flask  and  then  be  dissolved  out  with  water. 


AMMONIUM    BROMIDE  37 

sive  portions  of  10  cc.  of  hydrochloric  acid  solution  of  1.12 
sp.  gr.  until  the  washings  show  no  further  test  for  sulphates. 
(See  Note  5  (#).)  Then  transfer  the  crystals  to  a  porcelain 
dish  and  heat  gently,  while  stirring,  until  all  decrepitation 
ceases. 

Questions 

1 .  Why  must  the  hydrochloric  acid  gas  be  passed  through 
a  washing  bottle  ?     Why  is  the  safety  tube  necessary  ? 

2.  Why,  in  the  light  of  the  Mass  Law,  should  one  expect 
the  solubility  of  sodium  chloride  to  be  lessened  by  the  pres- 
ence of  hydrochloric  acid  ?     [It  may  be  stated  that  another 
effect  also  comes  into  play  here  which  likewise  tends  to 
lessen  the  solubility  of  sodium  chloride.] 

3.  Mention  two  possible  causes  for  the  very  consider- 
able amount  of  heat  produced  when  the  hydrochloric  acid 
gas  is  absorbed  by  the  solution  in  the  beaker. 

5.     AMMONIUM    BROMIDE 

Ammonium  bromide  could  be  prepared  by  the  neutral- 
ization of  ammonium  hydroxide  with  hydrobromic  acid, 

NH4OH  +  HBr  =  NH4Br  -f-  H2O. 

Since,  however,  hydrobromic  acid  is  a  more  expensive 
material  than  uncombined  bromine,  the  latter  would  have 
the  preference  as  a  source  of  bromine,  provided  it  yielded 
as  satisfactory  a  product.  Chlorine  or  bromine  reacts  as 
follows  upon  a  cold  solution  of  sodium  hydroxide,  as,  for 
example,  in  the  manufacture  of  bleaching  liquors, 

Br2  +  2NaOH  =  NaBr  +  NaBrO  +  H2O, 

with  the  formation  of  sodium  hypochlorite  or  hypobromite. 
Sodium  hypobromite  reacts  with  ammonium  hydroxide  ac- 
cording to  the  equation, 

3NaBrO  +  2NH4OH  =  3NaBr  +  5H2O  +  N2. 


38      ALKALI  AND  ALKALINE  EARTH  METALS 

Thus  the  action  of  bromine  upon  ammonium  hydroxide 
yields  only  ammonium  bromide  and  nitrogen  gas,  because 
even  if  the  primary  effect  were  to  yield  bromide  and 
hypobromite  in  equal  quantities,  as  is  the  case  if  sodium 
hydroxide  is  used,  the  ammonium  hypobromite  would  imme- 
diately react  with  fresh  ammonia  in  the  same  manner  as 
does  sodium  hypobromite. 

Procedure.  —  Place  55  cc.  of  concentrated  ammonia 
(sp.  gr.  0.90),  together  with  50  cc.  of  water,  in  a  flask,  which 
should  be  set  in  a  pan  of  ice  water.  Put  15.8  cc.  of  bromine 
in  a  small  separatory  funnel,1  and  add  it  a  drop  at  a  time  to 
the  ammonia,  rotating  the  flask  after  each  drop  until  the 
yellow  color  produced  by  the  bromine  has  completely  disap- 
peared. Do  not  allow  the  contents  of  the  flask  to  become 
heated  at  any  time,  as  a  dangerously  explosive  compound 
might  in  that  case  be  formed.  As  soon  as  a  permanent  yel- 
low color  is  produced,  stop  adding  bromine  and  add  at  once 
a  few  drops  of  ammonia  until  the  solution  has  again  become 
colorless.  Place  the  solution  in  an  evaporating  dish  on  top 
of  a  beaker  of  boiling  water,  and  let  the  salt  crystallize  as 
the  water  evaporates.  When  only  a  little  liquid  remains, 
separate  the  crystals  from  it  in  a  funnel,  and  dry  them  on 
a  porcelain  plate. 

Questions 

1.  What  products  would  be  formed  if   bromine  were 
added  to   a  solution  of   sodium  hydroxide  instead  of   am- 
monium  hydroxide,    (i)    if    the   solution   were   kept   cold  ? 
(2)  if  it  were  heated? 

2.  Experiment.  —  Add   about   10  drops  of  bromine  to 

xThe  stopcock  of  the  funnel  should  first  be  lubricated  with 
vaseline  or  grease  and  then  fastened  with  a  rubber  band.  Bromine 
produces  very  bad  burns  when  it  gets  upon  the  hands.  To  avoid 
danger  of  accident  ask  an  instructor  to  approve  the  apparatus  before 
beginning  actual  operations  with  the  bromine. 


STRONTIUM    HYDROXIDE  39 

10  cc.  of  a  cold  10  per  cent,  sodium  hydroxide  solution. 
Add  this  gradually  to  a  solution  of  ammonium  hydroxide, 
made  by  diluting  i  cc.  of  desk  reagent  with  10  cc.  of 
water.  Determine  what  gas  is  given  off. 

3.  Write   the  equation   showing  the  complete  reaction 
between    bromine    and    ammonia.       What   fraction    of    the 
entire  amount  of  ammonia  used  is  lost  through  formation 
of  nitrogen  gas  ? 

4.  Why   cannot   hydrobromic    acid    be    prepared    from 
potassium  bromide   by  a  method   analogous   to    that    used 
in  the  manufacture  of  hydrochloric  acid  ? 

5.  Explain  why,  from  the  standpoint  of  economy,  the 
method   of   preparation   above   outlined    is    superior  to  the 
direct  neutralization  of  ammonia  with  hydrobromic  acid. 


6.     STRONTIUM    HYDROXIDE    FROM    STRONTIUM 
SULPHATE 

One  of  the  most  important  sources  of  strontium  is  the 
mineral  celestite,  SrSO4.  By  reduction  with  charcoal  this 
can  be  converted  into  strontium  sulphide, 

SrSO4  +  4C  =  SrS  +  4CO, 

and  the  strontium  sulphide  by  treatment  with  copper  oxide 
and  water  can  be  made  to  yield  strontium  hydroxide, 

SrS  +  CuO  +  H2O  =  Sr(OH)2  +  CuS. 

Copper  oxide  is  in  the  ordinary  sense  insoluble ;  neverthe- 
less in  contact  with  water  it  does  yield  to  an  infinitesimal 
extent,  first  copper  hydroxide,  and  then  Cu++-ions, 

CuO  -]-  H2O  ^  Cu(OH)2  ^  Cu++  +  2OH-. 

Therefore,  since  copper  sulphide  is  a  far  more  insoluble 
substance  than  copper  oxide,  it  follows  that  the  few 


4O      ALKALI  AND  ALKALINE  EARTH  METALS 

Cu++-ions  from  the  latter  unite  with  the  S~~-ions  from  the 
strontium  sulphide  to  form  copper  sulphide,  which  pre- 
cipitates continuously,  while  the  copper  oxide  continuously 
goes  into  solution  to  resupply  Cu++-ions,  and  this  action 
continues  until  either  the  copper  oxide  or  the  strontium 
sulphide  is  exhausted. 

Strontium  hydroxide  crystallizes  with  8  molecules  of 
water,  Sr(OH)2.8H2O.  It  is  very  soluble  in  hot  water,  but 
sparingly  soluble  in  cold  water. 

Procedure.  —  Grind  50  grams  of  powdered  celestite  in 
a  porcelain  mortar  until  no  more  grit  is  felt  under  the 
pestle.  Add  20  grams  of  powdered  charcoal  and  continue 
to  grind  with  the  pestle  until  the  two  are  thoroughly  mixed. 
Place  the  mixture  in  a  clay  crucible,  pack  it  firmly,  and 
cover  it  with  a  layer  of  powdered  charcoal  \  inch  deep. 
Cover  the  crucible  with  a  close-fitting  cover  and  heat  it  in 
a  gas  furnace  for  one  hour,  at  a  bright  red  heat  (Instruc- 
tions as  to  regulating  the  flame).  After  the  contents  of  the 
crucible  have  cooled,  remove  the  layer  of  charcoal  from 
the  surface  and  bring  the  remainder,  after  crushing  it  to 
a  powder,  into  an  8-inch  porcelain  dish ;  add  300  cc.  of 
water,  bring  the  mixture  to  a  boil,  and  while  it  is  boiling 
add  copper  oxide,  a  little  at  a  time,  until  all  of  the  soluble 
sulphide  has  interacted  with  it,  —  about  40  grams  in  all. 
So  long  as  any  unchanged  strontium  sulphide  is  present 
the  solution  will  show  a  yellow  color,  which  may  be  observed 
by  letting  the  black  solid  settle  for  a  moment,  and  then 
looking  through  the  upper  layers  of  the  clear  liquid  at  the 
background  of  the  white  porcelain  dish.  As  soon  as  the 
yellow  color  has  entirely  disappeared,  the  strontium  sulphide 
has  all  reacted.  Crystals  of  strontium  hydroxide  separate 
rapidly  from  this  solution  when  it  cools.  Hence  it  must 
be  filtered  quickly  in  order  to  avoid  having  the  crystals 
form  in  the  filter  and  clog  it  completely.  Heat  50  cc.  of 


STRONTIUM    HYDROXIDE  41 

water  to  boiling  in  a  beaker,  and  keep  it  at  this  temperature 
until  it  is  .required.  Add  hot  water  to  the  dish  to  replace 
any  lost  by  evaporation,  and  pour  (Note  2)  the  hot  solution 
through  a  Witt  filter  (Note  4  (£)),  allowing  the  main  part  of 
the  residue  to  remain  in  the  dish.  Add  the  50  cc.  of  hot 
water  to  this  residue,  stir  it  thoroughly,  heating  it  for  a 
moment  over  the  flame,  and  then  pour  solution  and  residue 
into  the  filter  and  drain  out  all  of  the  liquid.  Transfer  the 
solution  to  an  Erlenmeyer  flask  (or  let  it  remain  in  the  filter 
bottle),  stopper  the  flask  to  exclude  the  air,  and  wrap  it 
with  a  towel,  so  that  the  solution -may  cool  slowly  and  larger 
crystals  may  be  formed.  Finally,  after  several  hours  cool 
the  solution  with  running  tap  water  and  then  collect  the 
crystals  on  a  Witt  filter.  Drain  the  crystals  for  a  moment, 
but  do  not  draw  too  much  air  through  them,  as  they  retain 
all  the  carbon  dioxide  it  contains.  Spread  the  moist  product 
on  filter  paper,  and  allow  it  to  dry  as  quickly  as  possible  by 
contact  with  the  air.  Stopper  it  in  a  sample  bottle  or  tube 
as  soon  as  it  is  dry. 

Questions 

1.  What   constituent  of   the  atmosphere   must   be  ex- 
cluded from  the  solution  while  crystallizing  and   as  much 
as  possible  from  the   crystals  while   drying?     How  would 
it  contaminate  the  preparation? 

2.  A  sample  of  the  preparation  should  dissolve  nearly 
clear  in  hot  water.     What  will  surely  cause  a  slight  cloudi- 
ness? 

3.  How   could    strontium    chloride   be   prepared   from 
strontium  sulphide  ? 

4.  Give  some  other  method  by  which  strontium  hydrox- 
ide could  be  obtained  from  strontium  sulphide  without  the 
use  of  copper  oxide. 

5.  Starting  with  the  mineral  strontium  carbonate,  how 


42      ALKALI  AND  ALKALINE  EARTH  METALS 

might  strontium  hydroxide  be  prepared  ?     Strontium  oxide  ? 
Strontium  chloride  ? 


7.  STRONTIUM  CHLORIDE  FROM  STRONTIUM 
SULPHATE 

Strontium  chloride  might  be  prepared  by  treating 
strontium  sulphide,  the  intermediate  product  in  the  last 
preparation,  with  hydrochloric  acid.  For  the  sake  of  illus- 
trating another  method,  however,  a  process  which  does  not 
require  the  use  of  a  furnace  is  here  employed  for  decom- 
posing the  strontium  sulphate. 

The  method  consists  in  first  converting  the  sulphate 
into  the  carbonate  by  boiling  it  with  a  concentrated  solution 
of  sodium  carbonate,  and  then  of  dissolving  the  carbonate 
in  hydrochloric  acid,  thereby  yielding  a  solution  of  the 
chloride.  The  conversion  of  solid  strontium  sulphate  into 
solid  strontium  carbonate  furnishes  an  interesting  illustration 
of  the  principle  of  mass  action,  for  the  solubility  of  these 
two  salts  in  pure  water  is  as  follows : 

Solubility  in  grams       Solubility  in  mols 
per  liter  per  liter 

SrSO4  o.on  0.0006 

SrCO3  o.oon  0*00007 

Strontium  sulphate  would  dissolve  in  the  solution  of  sodium 
carbonate  in  the  same  manner  as  it  would  in  pure  water 
until  it  had  saturated  the  solution  and  its  solubility  product, 
which  is  equal  to  0.0006  X  0.0006,  was  reached,  but  for 
the  fact  that  long  before  this  could  occur  the  solution 
would  be  supersaturated  with  respect  to  strontium  carbon- 
ate, whose  solubility  product  is  only  equal  to  0.00007  X 
0.00007.  Thus  strontium  carbonate  is  precipitated  con- 
tinuously as  strontium  sulphate  dissolves ;  and  since  the 


STRONTIUM     CHLORIDE  43 

solution  cannot  become  saturated  with  the  latter  so  long 
as  there  is  a  large  excess  of  carbonate  ions  present,  the 
solid  salt  finally  remaining  will  consist  entirely  of  strontium 
carbonate,  provided  a  sufficient  amount  of  sodium  carbon- 
ate were  employed.  The  reaction  which  takes  place  is, 
however,  reversible, 

SrSO4  +  Na2CO3  ^  SrCO3  -f  Na2SO4, 

and,  if  strontium  carbonate  were  boiled  with  a  solution  of 
sodium  sulphate,  the  solid  would  be  converted  into  the 
sulphate.  It  is  easy  to  deduce  that  if  the  ratio  of  the  con- 
centration of  the  ions  in  solution — — ^ —  is  greater 

cone  SO4~~ 

than  TL,  solid  strontium  sulphate  will  be  converted  into 
solid  carbonate. 

Procedure.  —  Take  50  grams  of  powdered  celestite. 
Grind  it  in  a  mortar  until  it  is  so  fine  that  it  no  longer 
feels  gritty  under  the  pestle.  Cover  it  in  an  8-inch  dish 
with  300  cc.  of  water,  add  60  grams  of  anhydrous  sodium 
carbonate,  and  boil  the  mixture  for  30  minutes,  stirring  it 
constantly  at  first.  Transfer  the  solution  and  solid  to  a 
tall,  narrow  beaker,  using  100  cc.  of  fresh  water  in  rinsing 
out  the  last  of  the  residue,  and  let  the  solid  matter  settle 
for  5  minutes.  Decant  off  the  liquid,  which  is  still  some- 
what cloudy,  but  from  which  the  essential  part  of  the  solid 
has  settled,  and  wash  the  residue  three  times  by  decantation 
with  400-500  cc.  of  water  (see  Note  5  (#),  first  paragraph). 
The  residue  is  now  sufficiently  free  from  soluble  sodium 
sulphate.  Transfer  about  y1^  of  the  moist  strontium  car- 
bonate to  another  beaker,  to  be  used  in  a  later  part  of  the 
process.  To  the  remaining  T9<y  add  50  cc.  of  hot  water,  and 
then  add  hydrochloric  acid,  drop  by  drop,  while  keeping 
the  mixture  at  the  boiling  temperature,  until  the  further 
addition  of  a  drop  of  acid  produces  no  more  effervescence. 


44      ALKALI  AND  ALKALINE  EARTH  METALS 

This  solution  now  contains  a  slight  excess  of  acid,  which 
is  to  be  neutralized  by  adding  the  remaining  TL  of  the 
strontium  carbonate.  Add  this,1  and  boil  the  mixture  for 
5  minutes.  The  solution  should  now  be  perfectly  neutral 
to  litmus.  If  it  is  acid,  it  shows  that  the  hydrochloric 
acid  was  added  carelessly  and  that  there  was  thus  more 
than  could  be  neutralized  by  the  strontium  carbonate. 
Filter  the  perfectly  neutral  solution,  and  evaporate  the 
filtrate  until  a  faint  scum  forms  on  removing  the  solution 
from  the  flame  and  blowing  vigorously  across  the  surface. 
Allow  the  solution  to  cool,  but  stir  occasionally  in  order  to 
obtain  a  uniform  crystal  meal  rather  than  a  cake  of  crystals. 
Finally,  drain  the  crystals  on  a  Witt  filter  (Note  4  (&) ;  evap- 
orate the  mother  liquor  to  crystallization  exactly  as  at  first, 
and  if  the  second  crop  of  crystals  is  pure  white,  add  it  to 
the  first  crop.  Dry  the  crystals  of  SrCl2.6H2O  by  spreading 
them  on  an  unglazed  plate  and  exposing  them  to  the  air 
at  room  temperature.  As  soon  as  they  are  dried  sufficiently 
so  as  not  to  cling  together,  bottle  them  at  once,  because 
they  are  somewhat  efflorescent. 

Questions 

1.  Explain   why   strontium    carbonate,   which    is    less 
soluble  in  pure  water  than  strontium  sulphate,  should  dis- 
solve readily  in  dilute  acids,  while  the  latter  salt  will  dissolve 
scarcely  any  more  in  acids  than  in  pure  water. 

2.  If  a  small  quantity  of  a  solution  of  strontium  chlo- 
ride were  added  to  a  solution  containing  equi-molal  quantities 
of  sodium  carbonate  and  sodium  sulphate,  what  would  be 
the  precipitate  formed  ? 

llf  any  iron  salt  is  present,  it  should  be  oxidized  by  adding  a  few 
drops  of  chlorine  water  and  boiling  a  moment  before  adding  the  stron- 
tium carbonate.  The  reason  for  such  a  procedure  will  be  made  clear 
later  under  Preparation  38,  Manganous  Chloride. 


BARIUM     OXIDE  45 

8.     BARIUM    OXIDE    AND    BARIUM    HYDROXIDE 
FROM    BARIUM    CARBONATE 

The  commercial  method  of  preparing  calcium  oxide 
(quicklime)  consists  in  heating  calcium  carbonate  (lime- 
stone) in  lime  kilns.  Barium  oxide  might  be  made  from 
barium  carbonate  according  to  the  same  principle,  except 
for  the  fact  that  the  temperature  required  for  the  decom- 
position of  barium  carbonate  is  so  high  as  to  make  such  a 
method  almost  impracticable.  This  greater  stability  of  the 
barium  salt  is  an  illustration  of  the  fact  that  barium  oxide 
is  even  more  strongly  basic  than  calcium  oxide.  The  reac- 
tion, BaCO3  ^  BaO  -|-  CO2,  'ls  to  some  extent  a  reversible 
reaction,  and  in  common  with  other  reversible  reactions  it 
may  be  made  to  progress  in  one  direction  or  the  other  by 
suitably  altering  the  concentration  of  the  substances  present 
in  the  reacting  mass.  Of  the  three  substances  involved  in 
this  reaction  the  only  one  which  can  be  removed  during  the 
course  of  the  reaction  is  the  carbon  dioxide.  It  is,  how- 
ever, not  enough  to  let  it  merely  pass  off  as  a  gas,  because 
there  is  already  present  in  the  atmosphere  in  the  furnace 
enough  carbon  dioxide  to  force  the  reaction  towards  the 
left.  The  carbon  dioxide  must  be  chemically  removed. 
This  is  accomplished  by  mixing  powdered  charcoal  with 
the  barium  carbonate,  for  carbon  reacts  with  carbon  diox- 
ide at  a  white  heat  and  gives  carbon  monoxide.  In  the 
following  procedure,  in  addition  to  the  charcoal,  a  little 
rosin  is  mixed  with  the  charge.  On  heating,  the  rosin 
decomposes  and  a  deposit  of  soot  is  formed,  which  in  this 
way  becomes  very  intimately  mixed  with  the  charge. 

The  barium  oxide  obtained  in  this  way  is  not  pure,  but 
contains  particles  of  charcoal  as  well  as  impurities  coming 
from  the  mineral.  It  is,  however,  very  suitable  for  the 
manufacture  of  barium  hydroxide,  into  which  it  is  con- 


46      ALKALI  AND  ALKALINE  EARTH  METALS 

verted  by  treatment  with  water.  Barium  hydroxide  is  ex- 
tremely soluble  in  hot  water,  but  sparingly  so  in  cold 
water,  from  which  it  separates  in  flake-like  crystals  of  the 
composition  Ba(OH)2.8H2O. 

Procedure.  —  Mix  100  grams  of  finely  powdered  barium 
carbonate  (use  either  the  mineral  witherite  or  the  artificially 
prepared  barium  carbonate)  with  10  grams  of  powdered 
charcoal  and  5  grams  of  rosin.  After  mixing  the  whole 
mass  very  thoroughly  in  a  mortar,  place  it  in  a  graphite 
crucible,1  press  it  down  firmly,  and  cover  it  with  a  layer 
of  charcoal  at  least  \  inch  deep.  Place  a  well-fitting  covei 
on  the  crucible  and  heat  the  whole  for  one  hour  to  as  high 
a  temperature  as  possible  in  the  gas  furnace.  After  the 
crucible  has  cooled,  remove  the  top  layer  of  charcoal,  break 
up  the  mass  of  barium  oxide,  and  add  it  rather  cautiously 
to  400  cc.  of  water  in  a  porcelain  dish.  (Note  whether  any 
heat  is  produced  when  barium  oxide  comes  in  contact  with 
water.)  Heat  the  mixture  in  the  dish  to  boiling ;  pour  the 
solution  through  a  large,  ordinary  filter  (Note  4  (^)),  letting 
the  clear  liquid  run  directly  into  a  500  cc.  flask.  Rinse  the 
residue  in  the  dish  with  75  cc.  more  of  boiling  water,  and 
pour  this  upon  the  filter  after  the  first  portion  has  nearly 
all  run  through.  Stopper  the  flask  and  allow  the  solution 
to  cool  slowly  to  room  temperature;  finally,  cool  it  nearly 
or  quite  to  o° ;  collect  the  crystals  on  a  Witt  filter,  and 
drain  and  dry  them  as  directed  for  strontium  hydroxide. 

Questions 

i.  How  does  barium  hydroxide  become  contaminated 
by  exposure  to  the  air  ?  Why  might  this  product  be  dried 

1  If  a  clay  crucible  is  used  it  should  be  lined  with  paper  before 
putting  in  the  charge.  The  paper  chars  and  leaves  a  layer  of  carbon 
which  separates  the  siliceous  material  of  the  crucible  from  the  basic 
barium  oxide  and  prevents  their  reacting  to  form  a  fusible  compound 


GENERAL    QUESTIONS    I  47 

with  less  contamination  by  exposure  to  the  air  out  of  doors 
than  to  the  air  of  the  laboratory  ? 

2.  The  mineral  witherite  often  contains  barium  sulphate 
as  an  impurity.     State  what  changes  this  substance  would 
undergo  during  the  above  process,  and  what  resulting  sub- 
stance   would    be   present   to    a   small   extent    in   the  final 
crystallized  product  and  to  a  greater  extent  in  the  mother 
liquor.     (Compare  preparation  of  strontium  hydroxide  from 
strontium  sulphate.)     Test  both  crystals  and  mother  liquor 
for   this    substance.      (How  ?)     The    final    product   can    be 
quite  satisfactorily  purified  from  it  by  one  or  two  recrystal- 
lizations.      How  x  might   it   be    removed    chemically?     (See 
Strontium  Hydroxide.) 

3.  Devise   a  method  for   preparing   barium   hydroxide 
from  barium  carbonate  by  which  the  use  of  a  furnace  may 
be  avoided.     Suggestion.  —  Make   use  of  the   difference   in 
solubility  of  barium  chloride  and  barium  hydroxide. 


GENERAL    QUESTIONS.    .1 

ALKALI   AND   ALKALINE   EARTH    METALS 

Experimen  ts 

(The  results  observed  are  to  be  recorded  in  the  laboratory  note- 
book at  the  time  the  experiments  are  performed.) 

i.  Place  a  few  small  lumps  of  marble  (pure  calcium 
carbonate)  in  a  small  iron  crucible,  or  a  porcelain  crucible 
lined  with  paper.  Cover  the  crucible  in  order  to  keep  in 
the  heat,  and  heat  it  quite  strongly  for  20  minutes  with  a 
Bunsen  flame.  When  the  product  has  cooled,  wet  each 
lump  with  a  single  drop  or  two  of  water  and  wait  a  few 
minutes,  if  necessary,  to  observe  the  effect.  Then  wet  the 
product  with  somewhat  more  water,  and  test  the  reaction 
of  the  moist  mass  towards  litmus. 


48      ALKALI  AND  ALKALINE  EARTH  METALS 

2.  Burn  a  strip  of   magnesium  ribbon,  held  with  iron 
pincers,  and  let  the  ash  fall  in  a  porcelain  dish.     Wet  the 
magnesium  oxide  with  a  single  drop  of  water  and  place  the 
moist  mass  on  a  strip  of  red  litmus  paper.     Note  the  rapidity 
and  intensity  with  which  the  litmus  is  turned  blue. 

3.  To  some  magnesium  chloride  solution,  add  (a)  some 
ammonium   hydroxide ;    (b)  some  ammonium   chloride    and 
then   some  ammonium   hydroxide.     Observe    in   each  case 
whether  magnesium  hydroxide  is  precipitated. 

4.  Dip  a  clean  platinum  wire  in  solutions  of   such  of 
the  chlorides  of  the  alkali  and  alkaline  earth  metals  as  are 
at   hand,    and  observe  the  color  imparted    to   the   Bunsen 
flame  when  the  wire  is  inserted  into  the  lower  part  of  the 
flame. 

Questions 

1 .  What  are  the  metals  of  the  alkali  group  ?     Of  the 
alkaline  earth  group  ?     Give  the  symbols  of  the  hydroxides 
and  chlorides  of  the  metals  of  each  group.     What  is  the 
characteristic  valence  of  each  group  ?     With  which  group 
is  the  ammonium  radical  NH4  classed  ? 

2.  Any  oxy-salt,  such  as  CaCO3  (=  CaO.CC^),  can  be  broken  up 
by  a  sufficiently  high  heat  into  a  basic  oxide  and  an  acid  oxide — for 
example,    CaO.CO2  -^  CaO  +  CO2.      The    higher    the   temperature 
necessary   to   accomplish    this,   the   greater   is    the   chemical    affinity 
between  the  two  oxides,  that  is,  the  more  strongly  basic  and  acidic, 
respectively,  are  these  two  components ;  and  therefore  in  a  series  of 
salts,   all   containing   the   same   acidic   oxide  —  for   example,    CaCOs, 
SrCO3,   BaCO3  —  the  greater   the  stability  of   the   salt,   the  stronger 
is  the  basic  oxide. 

Look  up  in  the  text-books  how  easily  the  carbonates  of 
sodium,  potassium,  calcium,  and  barium  are  decomposed  by 
heat  (compare  Experiment  i  and  Preparation  8);  also  to 
what  extent  and  how  readily  the  nitrates  of  the  same  metals 
are  decomposed.  From  these  data  arrange  the  oxides  of 
these  four  metals  in  the  order  of  their  strength  as  bases. 


GENERAL    QUESTIONS    I  49 

Which  are  more  strongly  basic,  the  oxides  of  the  alkali  or 
of  the  alkaline  earth  metals  ? 

3.  To  what  extent  are  the  hydroxides  of  sodium,  potas- 
sium,   calcium,    and    barium   dissociated   electrolytically    in 
aqueous    solution  ?     Does   there   seem   to   be    any   parallel 
between  the  basic  strength  of  an  oxide,  as  judged  from  the 
standard   explained   under  Question   2,   and   the  degree  of 
dissociation   of   the  base   derived  from   it  when    the   latter 
is  dissolved  in  water? 

4.  Compare  the  solubility  of  calcium  and   magnesium 
hydroxides    (Experiments    i    and    2).     Explain    the    results 
observed  in  Experiment  3. 

5.  Tabulate  the  colors  obtained  by  the  flame  tests  for 
the  alkali  and  alkaline  earth  metals  (see  Experiment  4  and 
text-books). 


CHAPTER    II 

ELEMENTS    OF    THE    THIRD    GROUP    OF    THE 
PERIODIC    SYSTEM 

Boron  and  aluminum,  the  first  two  members  of  this 
group,  are  the  only  ones  which  are  classed  among  the 
common  elements.  On  this  account,  and  also  because  the 
difference  in  properties  between  Family  A  and  Family  B 
is  far  less  marked  than  in  Groups  I  and  II,  the  whole 
group  is  taken  up  under  one  heading. 

The  characteristics  of  this  group  are  that  the  elements 
possess  a  valence  of  three,  and  that  the  oxides,  M2O3,  have 
but  a  weakly  developed  basic  character.  Boron,  in  fact, 
shows  practically  no  base-forming  properties,  but  forms 
rather  the  weak  boric  acid.  The  oxide  of  aluminum  dis- 
plays both  basic  and  acidic  properties ;  that  is,  it  is  am- 
photeric.  The  remaining  elements  are  more  distinctly 
base-forming  than  aluminum,  without,  however,  approaching 
in  any  way  the  alkaline  earth  metals  in  this  respect. 


9.     BORIC   ACID 

In  this  preparation,  borax,  the  sodium  salt  of  tetraboric 
acid,  is  chosen  as  the  source  of  boron.  Although  boron 
is  decidedly  a  non-metal,  still  its  acid-forming  characteristics 
are  not  highly  developed.  Its  acids,  therefore,  having  a  low 
degree  of  ionization,  are  readily  displaced  by  strong  acids 
from  solutions  of  their  salts.  Thus  tetraboric  acid  would  be 
set  free  from  borax  by  hydrochloric  acid;  but  since  this 
is  capable  in  aqueous  solution  of  combining  with  water,  and 
since  normal  boric  acid  (H3BO3)  is  but  slightly  soluble,  it 
is  this  substance  which  crystallizes  from  the  solution. 

Procedure.  —  Dissolve  100  grams  of  borax  in  300  cc.  of 
boiling  water.  Add  hydrochloric  acid  (about  41  cc.  of  cone, 
acid  1.2  sp.  gr.)  to  the  hot  solution  until  a  strip  of  blue 
litmus  paper  is  colored  bright  red.  Allow  to  cool,  when 
normal  boric  acid  will  crystallize  out.  Filter  off  the  crys- 
tals, using  a  Witt  filter,  and  purify  them  by  recrystallization, 
dissolving  in  300  cc.  of  hot  water.  Allow  to  cool  slowly, 
shaking  or  stirring  occasionally  to  prevent  the  crystals  from 
caking  together. 

Questions 

1.  Explain   the  relations   between   normal   boric   acid, 
metaboric  acid,  tetraboric  acid,  and  boric  anhydride.     Ex- 
periment. —  Place  a  few  grams  of  boric  acid  on  a  watch 
glass  upon  the  steam  table  (100-110°)  and  leave  for  \  hour. 
What  is  formed  ?     What  would  be  formed  if  the  acid  were 
heated  to  140°?     Suspend  a  little  of  the  acid  in  a  loop  of 
platinum  wire,   and   heat  in   the   Bunsen.  flame.     What   is 
formed  ? 

2.  Experiment.  —  Place  a  few  grains  of  boric  acid  in 
a  small  porcelain  dish,  cover  it  with  5  cc.  of  alcohol,  set  fire 

53 


54         ELEMENTS  OF  THE  THIRD  GROUP 

to  it,  and  observe  the  color  of  the  edges  of  the  flame,  espe- 
cially when  stirring  and  when  the  alcohol  is  almost  burned 
out.  Repeat,  using  borax  instead  of  the  boric  acid,  and 
again,  using  borax  moistened  with  concentrated  sulphuric  acid. 

What  causes  the  green  color  of  the  flame,  and  why  is 
it  not  observed  with  borax  alone  ? 

Repeat  if  necessary  the  last  part  of  the  preceding 
experiment,  noticing  the  color  imparted  to  the  flame  while 
the  orthoboric  acid  is  first  melting,  and  again  when  a  clear 
bead  of  boric  anhydride  is  obtained. 

What  conclusions  can  you  make  from  these  experiments 
regarding  the  volatility  of  boric  acid  and  of  boric  anhydride  ? 

3.  What  effect  has  a  solution  of  borax  upon  litmus? 
Explain  what  is  thus  shown  regarding  the  strength  of  boric 
or  tetraboric  acid.     Explain  why  litmus  will  not  be  turned 
a  bright  red  until  more  than  two  molecules  of  HC1  have 
been  added  to  one  molecule  of  borax. 

4.  How  can   chloride   of   boron    be    prepared  ?      How 
does  this  -substance  behave  when  treated  with  water  ?     How 
would  it  behave  if  boron  were  a  strongly  metallic  element  ? 

10.    ALUM    FROM    CRYOLITE 
(BY-PRODUCT:  SODIUM  CARBONATE) 

A  characteristic  series  of  compounds  of  the  trivalent 
metals  are  the  alums,  of  which  the  potassium  aluminum,  or 
common  alum,  K2SO4.Al2(SO4)3.24H2O,  is  typical..  These 
compounds  are  particularly  interesting  from  the  readiness 
with  which  they  can  be  produced  in  large  and  beautiful 
crystals,  and  from  the  fact  that  any  of  the  univalent  alkali 
metals  may  take  the  place  of  potassium,  or  any  one  of  a 
large  number  of  the  trivalent  metals  may  take  the  place  of 
aluminum  in  the  common  alum  without  altering  the  form 
or  chemical  nature  of  the  crystals  produced. 


ALUM  55 

Of  the  many  raw  materials  which  might  serve  as  the 
source  of  aluminum,  cryolite,  the  double  fluoride  of  sodium 
and  aluminum,  3NaF.AlF3  or  Na3(AlF6),  has  been  chosen, 
partly  because  it  actually  serves  as  an  important  source  of 
aluminum  compounds,  and  partly  because  its  decomposition 
illustrates  important  chemical  reactions  and  manipulations. 

This  mineral,  although  itself  insoluble  in  water,  will,  if 
boiled  with  a  suspension  of  milk  of  lime  (Ca(OH)2),  undergo 
a  metathesis,  with  the  formation  of  insoluble  calcium  fluo- 
ride, and  a  soluble  salt  of  aluminum,  in  which  this  metal 
plays  the  part  of  an  acid-forming  element, 

3NaF.AlF3  +  3Ca(OH)2  =  3CaF2  +  Na3AlO3  +  3H2O. 

By  removing  the  insoluble  residue  left  by  this  reaction  from 
the  liquid,  a  separation  of  the  aluminum  from  the  fluorine  is 
accomplished ;  but  on  the  laboratory  scale  this  separation 
is  difficult  to  carry  out  on  account  of  the  colloidal  nature 
of  the  residue.  If  filtration  were  resorted  to,  the  pores  of 
the  filter  would  be  immediately  clogged  with  the  gelatinous 
precipitate,  so  that  the  liquid  would  run  so  slowly,  even 
with  suction,  that  an  undue  length  of  time  would  be  spent 
in  accomplishing  the  separation.  Therefore  the  method 
which  will  be  employed  is  that  of  sedimentation  after  stir- 
ring up  with  a  large  amount  of  water.  It  must  be  borne 
in  mind,  however,  that  this  device  is  adopted  only  to  meet 
the  requirements  of  laboratory  practice,  for  on  a  commer- 
cial scale  the  expense  of  evaporating  the  large  amount  of 
water  to  obtain  soda  crystals,  one  of  the  by-products  of  the 
process,  would  be  prohibitive.  The  separated  sludge,  con- 
sisting of  calcium  fluoride  and  all  excess  of  calcium  hydrox- 
ide, is  of  no  great  value,  and  will  be  discarded,  although 
it  could  be  used  as  a  source  of  fluorine  compounds ;  for 
example,  hydrofluosilicic  acid  (H2SiF6). 

From  the  clear  solution  of  sodium  aluminate  the  alumi- 


50          ELEMENTS  OF  THE  THIRD  GROUP 

num  is  precipitated  by  displacement  of  the  weak  aluminic 
acid  from  the  salt  by  the  action  of  the  stronger  carbonic 
acid, 

2Na3A103  +  3H20  +  3C02  =  3Na2CO3  -f  2H3A1O3. 

This  precipitated  aluminic  acid,  H3A1O3  (or,  as  more  fre- 
quently named,  aluminum  hydroxide,  A1(OH)8),  is  also  a 
very  gelatinous  substance,  and  can  likewise  only  be  sepa- 
rated, within  a  reasonable  length  of  time,  by  means  of 
sedimentation. 

The  aluminum  hydroxide  is  treated  with  the  calculated 
amount  of  sulphuric  acid  whereby  the  soluble  salt,  aluminum 
sulphate,  is  obtained.  To  this  solution  is  added  the  calcu- 
lated amount  of  potassium  sulphate,  and  then  the  alum  is 
allowed  to  crystallize. 

The  ideal  conditions  for  obtaining  large,  clear  crystals  — 
which  constitutes  the  beauty  of  this  as  a  laboratory  prepa- 
ration—  are,  that  a  solution  which  is  just  saturated  with  alum 
may  be  slowly  concentrated  by  spontaneous  evaporation  at 
a  nearly  constant  temperature.  Such  conditions  are  found 
in  industrial  works,  where  the  evaporation  of  the  solution 
in  large  vats  yields  beautiful  crystals,  often  of  enormous 
size ;  but  these  necessary  conditions  are  almost  impossible 
to  realize  in  a  small  laboratory  preparation,  and  another 
method  is  adopted  to  give  more  rapidly  and  more  surely 
the  desired  results. 


100  grams  water  dissolve  the  given  number  of  grams  of 
K2S04.A12(S04)3.24H2O. 


Temp.  .  . 

OP 

5° 

10° 

15° 

20° 

25° 

30° 

40° 

50° 

60° 

70° 

Grams  .  . 

5.6 

6.6 

7.6 

9.6 

11.4 

14.1 

16.6 

24 

36 

57 

110 

ALUM  57 

From  the  accompanying  table  it  is  seen  that  the  solubility 
increases  rapidly  with  the  temperature.  If  a  solution  is 
saturated  at  35°  and  then  cooled  to  15-20°  about  one-half 
of  the  alum  will  separate  out,  but  ordinarily  in  the  form  of 
a  mass  of  very  minute  crystals.  In  order  to  obtain  large 
crystals  during  the  cooling,  three  precautions  are  necessary : 

(1)  A  few  minute  crystals  must  be  added  to  serve  as  nuclei 
for  the  crystallization  before   setting  the   solution  to  cool. 

(2)  Dust  must  be  excluded,  since  dust  particles  might  serve 
as  nuclei  for  the  formation  of  a  great  number  of  minute 
crystals.     (3)  The  cooling  must  take   place  very  slowly  in 
order   that   the  crystal   faces    may  be    built   up   uniformly. 
This  can  be  accomplished  if  the  crystallizing  dish  is  insu- 
lated by  being  covered  with  a  glass  plate  and  wrapped  with 
a  towel.     If  the  solution  set  to  crystallize  is  saturated  at  a 
much  higher  temperature  than  35°,  it  will  be  found  that  the 
crystallization  will  proceed  so  rapidly  that  it  will  be  quite 
impossible,  even  if  the  above  precautions  are  all  observed, 
to  obtain  good  crystals. 

Procedure.  —  Slake  46  grams  of  quicklime  and  then  mix 
in  a  large  casserole  with  250  cc.  of  water  to  make  milk  of 
lime.  Stir  into  this  50  grams  of  finely  powdered  cryolite, 
and  heat  to  boiling  while  stirring  constantly.  Continue 
boiling  for  an  hour,  adding  water  to  replace  that  lost  by 
evaporation  and  stirring  sufficiently  to  avoid  spattering. 
At  the  end  of  that  time  add  a  part  of  \\  liters  of  boiling 
water  to  the  thick  mass  in  the  dish,  and  transfer  it,  together 
with  the  rest  of  the  hot  water,  to  a  tall,  wide-mouth,  2-liter 
bottle,  and  set  aside  to  settle  until  the  next  period.  The 
sludge  should  have  settled  so  as  to  occupy  not  more  than 
one-fifth  of  the  volume  of  the  liquid.  Siphon  off  as  much 
clear  liquid  as  possible  without  drawing  over  any  of  the 
precipitate.  Then  add  \\  liters  more  of  hot  water  to  the 
bottle,  stir  and  again  let  settle,  and  draw  off  the  clear 


58  ELEMENTS    OF    THE    THIRD    GROUP 

liquor.  The  residue  in  the  bottle  may  be  thrown  away. 
Combine  all  the  solution  in  a  large  bottle,  and  pass  in  car- 
bon dioxide  from  a  Kipp  generator  until  all  of  the  alumina 
is  precipitated.  Test  to  see  if  this  is  accomplished  at  the 
end  of  ^  hour  by  stopping  the  carbon  dioxide  stream  and 
letting  the  precipitate  settle  enough  to  pour  off  a  little 
clear  liquor  into  a  beaker.  Pass  carbon  dioxide  into  this 
for  a  few  minutes ;  if  no  precipitation  occurs  it  shows  that 
all  of  the  alumina  has  already  fallen  out  of  the  solution. 
If  a  precipitate  does  appear,  the  treatment  of  the  entire 
solution  with  carbon  dioxide  must  be  continued  until  all 
the  alumina  is  thrown  out.  Then  let  the  solution  settle 
until  the  precipitate  occupies  less  than  one-sixth  of  the 
entire  volume.  Siphon  off  the  clear  liquid  and  evaporate 
it  to  dryness  in  a  porcelain  dish.  Powder  and  preserve 
the  sodium  carbonate  so  obtained.  Stir  up  the  precipitate 
left  in  the  bottle  with  i^  liters  of  hot  water;  let  settle,  and 
siphon  off  and  discard  the  clear  liquid,  since  it  will  not 
contain  sufficient  sodium  carbonate  to  pay  for  its  evapora- 
tion. To  the  suspension  of  aluminum  hydroxide  left  in  the 
bottle  add  30  cc.  of  concentrated  sulphuric  acid,  and  warm, 
if  necessary,  to  effect  complete  solution.  Add  42  grams  of 
potassium  sulphate,  and  warm  until  dissolved.  The  solu- 
tion should  now  be  perfectly  clear ;  if  not,  filter.  If  the 
volume  exceeds  400  cc.  evaporate  to  that  bulk,  and  while 
still  hot  transfer  it  to  the  crystallizing  dish  (an  8-inch 
porcelain  dish  will  answer).  When  cooled  to  55-50°,  drop 
8  to  10  very  small  alum  crystals  into  the  solution,  cover 
immediately  with  a  glass  plate,  wrap  the  whole  in  a  towel, 
and  set  where  it  will  not  be  disturbed  until  the  next  exer- 
cise. Remove  the  few  large  crystals  formed  and  preserve 
them.  Evaporate  the  mother  liquor  not  quite  to  £  its  bulk 
(say  T%-),  and  set  this  to  crystallize  in  exactly  the  same 
manner  as  before.  Add  the  crystals  so  obtained  to  the 
first  lot. 


ALUMINUM    SULPHIDE  59 

Questions 

1.  Of  what   does  the   insoluble   residue   consist  which 
remains  after  boiling  the  cryolite  with  milk  of  lime  ?     When 
this  is  discarded   after  partial  washing  according  to  direc- 
tions, what  proportion  of  the  soluble  aluminum  salt  is  lost 
with  it  ?     (Note  that  the  actual  solid  material  of  the  slime 
in    question    occupies    an    inappreciable    volume    as    com- 
pared with  the  liquid  in  which  it  is  suspended,  even  after  the 
slime  has  been  settling  for  several  days.     See  Note  5  (b)  on 
page  13.) 

2.  State  what  ions  are  produced  by  aluminum  hydrox- 
ide (i)  when  it  acts  as  an  acid;   (2)  when  it  acts  as  a  base. 
Compare  its  strength  as  an  acid  and  as  a  base  with  that 
of  other  common  electrolytes. 

Define  an  amphoteric  electrolyte. 

3.  Experiment.  —  Add  a  solution  of  sodium  hydroxide, 
drop  by  drop,  to  5  cc.  of  a  solution  of  aluminum  chloride, 
continuing  until  a  considerable  amount  is  added  after  the 
first  effect  is  observed.     Preserve  the  contents  of  the  tube. 

Repeat,  using  a  solution  of  ammonium  hydroxide  instead 
of  one  of  sodium  hydroxide. 

To  the  tube  saved  from  the  first  part  of  this  experiment 
add  a  solution  of  ammonium  chloride. 

Show  how  the  difference  in  the  degree  of  dissociation  of 
the  two  bases  accounts  for  the  difference  in  the  first  two 
cases.  State  how  it  may  also  explain  the  result  in  the 
third  case. 

ii.     ALUMINUM    SULPHIDE 

This  compound  cannot  be  prepared  in  the  wet  way  for 
the  reason  that  it  is  decomposed  by  water.  It  is,  however, 
readily  prepared  by  the  action  of  metallic  aluminum  upon 
lead  sulphide  at  a  high  temperature. 


6O         ELEMENTS  OF  THE  THIRD  GROUP 

Procedure.  —  Mix  133  grams  of  finely  powdered  galena1 
(lead  sulphide  ore)  with  10  grams  of  granulated  aluminum. 
Place  in  a  small  uncovered  Hessian  crucible,  and  heat  as 
strongly  as  possible  in  a  gas  furnace  until  the  charge  in 
the  crucible  commences  to  react  and  a  bright  glow  extends 
throughout  the  whole  mass.  Leave  two  or  three  minutes 
longer  in  the  furnace  at  a  white  heat  in  order  that  the  re- 
action may  be  entirely  completed  and  that  the  charge  may 
become  quite  liquid  (a  thin  solid  crust  of  aluminum  oxide  may 
form  on  the  surface).  Then  pour  into  a  dry,2  clean  iron  pan  ; 
when  cool  detach  the  brittle  aluminum  sulphide 3  from  the 
lead  button  and  stopper  the  sulphide  tightly  in  a  test  tube. 
The  lead  may  be  put  in  a  box  provided  for  scrap  lead. 

Questions 

1.  Experiment.  —  Drop  a  small  lump  of  aluminum  sul- 
phide (at  the  hood)  into  a  test  tube  of  water.     What  is  the 
gas  formed  and  what  is  the  insoluble  residue  left  ?     What 
are  the  free  acid  and  base  into  which  the  salt  is  resolved 
by  the  action  of  water?     State  the  specific  nature  of  both 
the  acid  and  the  base,  which  makes  this  reaction  possible, 
instead  of  the  reverse  reaction  (neutralization),  which  is  the 
common  one  between  an  acid  and  a  base. 

2.  If  uncombined   sulphur  were  used   instead  of  lead 
sulphide    as    the    source    of    sulphur   for   this    preparation, 
suggest  reasons   why  the  process   would  be   more  difficult 
to  control. 

1 A  fairly  pure  sample  of  the  mineral  must  be  used  in  order  to 
obtain  good  results. 

2  If  the  iron  pan  should  chance  to  be  moist,  the  heat  of  the  charge 
would   cause    the   formation. of   steam   with    explosive   rapidity,   thus 
throwing  about  the  red-hot  material. 

3  The  aluminum  sulphide  must  be  handled  entirely  under  the  hood. 
None  of  it  must  be  allowed  to  fall  on  the  desks  or  floor  of  the  open 
laboratory,  because  it  reacts  with  the  moisture  of  the  air,  producing  a 
strong  odor  of  hydrogen  sulphide. 


GENERAL    QUESTIONS    II  6 1 

GENERAL    QUESTIONS.     II 

ELEMENTS   OF   THE   THIRD   GROUP    OF   THE    PERIODIC   SYSTEM 

Experiments 

(The  results  observed  are  to  be  recorded  in  the  laboratory  notebook 
at  the  time  the  experiments  are  performed.) 

1.  Evaporate  a  little  of  an  aluminum  nitrate  or  alumi- 
num chloride  solution,  and  heat  the  residue  carefully  over 
a  free  flame  until  it  is  just  dry.     Examine  the  residue  and 
test  its  solubility  in  water. 

2.  (a)  Add  ammonium  hydroxide,  drop  by  drop,  to  a 
solution  of  any  soluble  salt  of  aluminum  until  a  precipitate 
has  formed.     (#)  Then   add   a   considerable   excess  of   the 
reagent,     (c)  Collect  a  little  of  the  precipitate  of  aluminum 
hydroxide  on  a  filter,  wash  it  entirely  free  from  ammonia 
by  means  of  hot  water,  and  then  test  the  reaction  of  the 
moist  solid  towards  litmus. 

3.  Repeat  (a)  and  (fr)  of  Experiment  2,  using  sodium 
hydroxide    instead    of   ammonium    hydroxide.      Divide    the 
solution  resulting  from  (ft)  into  two  parts:  (c)  To  one  part 
add    hydrochloric   acid,    drop  by   drop,   until    a   precipitate 
forms,   then   add    an   excess,     (d)  To   the   other   part    add 
ammonium  chloride. 

Questions 

i.  Boron  and  aluminum  are  the  only  members  of  the 
third  group  of  the  periodic  system  of  the  elements  which 
are  of  common  occurrence  or  of  great  importance.  Turn  to 
a  table  of  the  periodic  system.  From  what  you  know  of  the 
regularities  of  this  system  what  should  you  predict  would 
be  the  valence  of  scandium,  lanthanum,  indium,  thallium  ? 
Write  the  symbols  of  the  sulphates  of  these  elements.  Find 
out  from  a  text-book  what  chlorides  of  thallium  are  known 
to.  exist.  Which  one  is  characteristic  of  the  position  of 


62          ELEMENTS  OF  THE  THIRD  GROUP 

thallium  in  the  third  family  ?  Is  it  of  common  occurrence 
for  an  element  to  display  different  valences?  Does  this 
occur  in  the  alkali  or  alkaline  earth  groups  ? 

2.  How  readily  is    aluminum    nitrate    decomposed    by 
heat  (Experiment  i)  ?     Compare  it  with  the  nitrates  of  the 
alkali    and    alkaline    earth    metals  (General    Questions,    I). 
What  is  to  be  said  of  the  stability  of  aluminum  carbonate  ? 
How  does  aluminum  compare  with  the  alkali  and  alkaline 
earth  metals  as  a  base-forming  element  ? 

3.  Compare    aluminum    hydroxide    with    sodium    and 
magnesium  hydroxides  as  regards  solubility  and  degree  of 
electrolytic  dissociation. 

4.  Explain  fully  the  behavior  of   aluminum  hydroxide 
as  an  amphoteric  substance  (see  Experiment  3). 

5'  In  the  third  group  of  the  periodic  system  the  elements  grow 
more  strongly  base-forming  as  the  atomic  weight  increases  and  more 
strongly  acid-forming  as  the  atomic  weight  grows  less.  Thus  with 
boron,  base-forming  properties  are  almost  lacking  and  acid-forming 
characteristics  are  more  strongly  developed  than  with  aluminum. 

How  can  anhydrous  boron  chloride  and  anhydrous 
aluminum  chloride  be  prepared  ?  How  do  each  of  these 
substances  behave  when  treated  with  water,  and  how  can 
the  difference  in  the  metallic  character  of  the  elements 
be  judged  from  the  difference  in  this  behavior  ? 


CHAPTER    III 

HEAVY    METALS   OF   THE   FIRST   TWO   GROUPS 
OF   THE    PERIODIC   SYSTEM 

The  metals  coming  under  this  heading  constitute  the 
right  hand  or  B  Families  in  Groups  I  and  II  of  the  periodic 
system.  They  possess  high  specific  gravities,  and  chemic- 
ally they  are  far  less  active  than  the  metals  of  the  corre- 
sponding A  Families, — they  being  not  greatly,  or  not  at  all, 
affected  by  the  atmosphere  or  by  water.  They  are  distinctly 
base-forming,  in  that  their  oxides  yield  fairly  stable  salts 
with  the  strong  acids ;  but  their  basic  properties  are  com- 
paratively weak,  and  the  oxides  of  some  of  them  show 
very  feeble  acidic  properties  as  well. 

Copper,  silver,  and  gold  in  Group  I  show  a  similarity 
to  sodium  and  potassium  principally  in  the  fact  that  they 
form  certain  compounds  of  the  same  type,  for  example, 
M2O  and  MCI.  Zinc,  cadmium,  and  mercury  in  Group  II 
resemble  calcium,  barium,  and  strontium  in  that  they  form 
compounds  of  the  types,  MO,  MSO4,  MC12,  etc.  In  other 
respects  the  divergence  in  the  properties  of  the  elements 
of  the  A  and  B  Families  is  at  a  maximum  in  these  two 
groups. 


12.     CRYSTALLIZED    COPPER    SULPHATE    (BLUE 
VITRIOL)    FROM    COPPER    TURNINGS 

On  account  of  the  fact  that  copper  has  not,  like  zinc, 
iron,  etc.,  the  power  of  displacing  hydrogen  from  acids,  it 
is  not  possible  to  dissolve  it  directly  in  dilute  sulphuric  acid. 
But  although  the  metal  itself  is  so  difficult  to  affect  with 
acids,  nevertheless  copper  oxide  is  readily  dissolved ;  and 
thus  the  problem  becomes  to  convert  copper  into  its  oxide. 
The  cheapest  source  of  oxygen  is  the  atmosphere,  and  on 
the  commercial  scale  the  usual  method  of  obtaining  copper 
sulphate  from  scrap  copper  is  to  allow  dilute  sulphuric  acid 
to  drip  slowly  over  the  latter,  to  which  is  given  free  access 
of  air.  Since,  however,  this  method  would  be  too  time- 
consuming  to  apply  on  a  small  scale,  nitric  acid  will  be 
employed  instead  of  air  as  the  oxidizing  agent. 

3Cu  +  2HNO3  =  3CuO  +  H2O  +  2NO. 
CuO  +  H2SO4  =  CuSO4  +  H2O. 

Procedure.  —  Heat  50  grams  of  copper  turnings  in  an 
iron  pan  until  all  oily  matter  is  burned  and  the  metal  has 
become  coated  with  oxide.  In  a  500  cc.  porcelain  dish 
treat  the  ignited  copper  turnings  with  300  cc.  of  dilute 
(1:4)  sulphuric  acid  and  125  cc.  of  dilute  nitric  acid 
(sp.  gr.  1.2).  Warm  in  the  hood  over  a  Bunsen  flame  for 
20  minutes ;  if  any  metallic  copper  remains  undissolved,  pour 
the  solution  off  from  it  and  dissolve  it  with  a  few  cubic 
centimeters  of  fresh  nitric  acid  and  twice  as  much  sulphuric 
acid.  If  the  solution  is  not  perfectly  clear,  filter  it  while 
still  at  the  boiling  temperature ;  then  cool  as  rapidly  as 
possible,  stirring  to  get  a  crystal  meal.  Separate  the  crys- 
tals from  the  mother  liquor  by  use  of  a  Witt  filter.  Evapo- 
rate the  mother  liquor  somewhat,  and  obtain  a  second  crop 

65 


66  HEAVY    METALS 

of  crystals,  throwing  away  the  mother  liquor  from  this 
crystallization.  Dissolve  all  of  the  crystals  by  heating  in 
their  own  weight  of  water;  set  the  solution  away  to  cool 
slowly,  and  after  several  hours  remove  the  crystals  which 
have  formed  and  evaporate  the  remaining  solution  to 
obtain  another  crop  of  crystals. 

Questions 

1.  Explain   why   copper   will    not    dissolve    in    dilute 
sulphuric  acid. 

2.  Write  the  equation  for  the  reaction  of  copper  with 
concentrated   sulphuric   acid.      Analyze   this    reaction   and 
show  in  what  manner  the  copper  is  oxidized. 

3.  The   quantities   of   sulphuric  and   nitric   acid  used 
were  considerably  in  excess  of  what  is  theoretically  neces- 
sary  to   react   with   the   given   amount   of    copper.      Give 
reasons  why  the  final  product  is  not  likely  to  be  contami- 
nated with  copper  nitrate  or  nitric  acid  or  with  free  sulphuric 
acid. 

4.  How  can  copper  sulphate  be  obtained  from  copper 
sulphide  on  a  commercial  scale  ? 

13.     CUPROUS    CHLORIDE 

Although  copper  itself  is  not  readily  oxidized,  yet  when 
oxidation  is  once  induced  it  proceeds  under  most  condi- 
tions at.  once  to  the  "  ic  "  state  (corresponding  to  the  oxide 
CuO).  This  is  in  accordance  with  the  fact  that  cuprous 
salts  are  oxidized  with  the  greatest  readiness,  so  that,  in 
the  presence  of  any  oxidizing  agent  powerful  enough  to  oxi- 
dize metallic  copper,  any  cuprous  salt  which  might  first  be 
formed  would  instantly  be  oxidized  to  cupric.  Therefore 
the  readiest  means  of  preparing  a  cuprous  salt  is  to  first 
prepare  the  cupric  salt  and  then  partially  reduce  it. 


CUPROUS    CHLORIDE  6/ 

If  metallic  copper  and  cupric  chloride  be  made  to  react, 
the  former  will  act  as  a  reducing  agent,  the  latter  as  an 
oxidizing  agent,  and  both  will  pass  into  the  "  ous "  state 
of  oxidation : 

Cu  +  CuCl2  =  2CuCl. 

Cuprous  chloride  is  soluble  in  a  concentrated  solution  of 
hydrochloric  acid  with  formation  of  a  complex  compound ; 
the  latter,  however,  is  broken  up  on  dilution  with  water,  and 
insoluble  cuprous  chloride  is  precipitated.  Cuprous  chloride 
is  so  readily  oxidized  that,  even  in  contact  with  moist  air  or 
suspended  in  water  containing  dissolved  air,  it  is  changed  to 
a  cupric  salt : 

2CuCl  +  2HC1  -f  O  =  H2O  +  2CuCl2, 
or  2CuCl  +  H2O  +  O  = 

Hence  a  great  deal  of  care  will  be  necessary  in  washing 
and  drying  the  cuprous  chloride  to  prevent  its  becoming 
discolored  in  consequence  of  oxidation.  When  perfectly 
dry,  however,  or  when  covered  with  dry  ether,  the  oxidation 
takes  place  very  much  more  slowly. 

.-  Procedure.  —  Dissolve  43  grams  of  crystallized  cupric 
chloride,  CuCl2.2H2O,  in  100  cc.  of  water  and  100  cc.  of 
concentrated  hydrochloric  acid  (1.2).  Place  the  solution, 
together  with  25  grams  of  fine  copper  wire  or  clean  copper 
turnings,  in  an  Erlenmeyer  flask ;  suspend  a  short-stemmed 
funnel  in  the  neck  of  the  flask  to  prevent  in  part  the  escape 
of  acid  vapors ;  raise  the  liquid  to  the  boiling  temperature, 
and  keep  it  at  that  point,  but  without  allowing  it  to  actually 
boil  appreciably,  until  the  green  color  has  disappeared.1 
If  then  a  few  drops  of  the  liquid  added  to  a  test  tube  of 
water  impart  no  blue  color  to  it,  the  reaction  is  complete. 

1  If  no  impurities  are  present  the  solution  becomes  colorless.  The 
color  is  more  likely,  however,  to  change  from  green  to  deep  brown. 


68  HEAVY    METALS 

Filter  the  solution  through  asbestos  (see  Note  4  (//)),  after 
first  moistening  the  asbestos  felt  with  concentrated  hydro- 
chloric acid.  Pour  the  clear  solution  into  a  tall,  common 
bottle  containing  2  liters  of  water.  Allow  the  white  pre- 
cipitate to  settle,  pour  off  the  clear  liquid,  and  stir  up  the 
precipitate  again  with  200  cc.  of  very  dilute  hydrochloric 
acid  (about  i  cc.  of  the  acid  of  1.2  sp.  gr.  to  100  cc.  of 
water).  After  the  precipitate  settles  again,  pour  off  the  clear 
liquid,  collect  the  precipitate  on  a  Witt  filter  (Note  4  (<£)), 
and  wash  it,  first  with  very  dilute  hydrochloric  acid,  then 
with  two  successive  portions  of  15  cc.  each  of  alcohol,  and 
then  with  two  portions  of  ether  of  like  volume  (Note  5  («)). 
During  the  above  treatment  the  precipitate  should  be  kept 
out  of  contact  with  the  air  as  much  as  possible,  by  keeping 
it  covered  with  liquid  all  the  time  until  it  is  freed  from 
water  by  means  of  the  alcohol  and  ether.  After  the  ether 
has  drained  off,  break  up  the  caked  cuprous  chloride  and 
leave  it  in  a  warm  and  dry  place  until  the  adhering  ether 
has  completely  evaporated.  Pulverize  the  product  and 
stopper  it  tightly  in  a  dry  test  tube.  The  product  should 
be  white.  If  it  is  discolored  it  must  be  dissolved  in  con- 
centrated hydrochloric  acid,  poured  into  water,  and  washed 
and  dried  as  before.  x 

Questions 

1.  Does    copper   resemble    silver    more   when    in    the 
cupric  or  the  cuprous  condition? 

2.  Experiment.  —  Expose  a  little  cuprous  chloride  cov- 
ered with  water  to  the  sunlight.     What  property  is  observed 
in  which  cuprous  chloride  is  like  silver  chloride  ? 

3.  Experiment.  —  Place  2  to  3  grams  of  cuprous  chlo- 
ride in  the  bottom  of  a  dry  test  tube.     Have  a  stopper  fitted 
to   the   tube,  then    fill   it   completely  with   ammonia   water 
(sp.  gr.  0.96)  and  immediately  stopper  it  tightly,  allowing  no 


AMMONIUM    AND    COPPER    SULPHATE  69 

air  bubble  to  remain  at  the  top.  Invert  the  tube  a  number 
of  times  until  the  salt  is  dissolved.  At  this  point  the  solu- 
tion should  be  nearly  colorless,  and  it  would  be  quite  so 
if  the  salt  had  been  pure  and  air  had  been  completely 
excluded.  Unstopper  the  tube  and  pour  quickly  its  con- 
tents into  a  bottle  of  about  200  cc.  capacity,  and  imme- 
diately cork  the  latter  air-tight.  Shake  the  bottle  vigorously 
for  3  minutes,  then  place  the  mouth  under  water  and  open 
the  cork.  From  its  quantity,  infer  what  gas  is  absorbed  out 
of  the  air  in  the  bottle.  What  is  the  change  in  the  condi- 
tion of  the  copper  salt  which  causes  its  change  in  color? 
Formulate  the  reaction. 


14.     AMMONIUM    AND    COPPER    SULPHATE 
(NH4)2SO4.CuSO4.6H2O 

One  example  of  the  formation  of  a  double  salt  has 
already  been  illustrated  in  the  preparation  of  alum,  in 
which  there  separates  from  a  solution  containing  the  two 
simple  salts  K2SO4  and  A12(SO4)3,  not  a  mixture  of  the 
crystals  of  each  salt,  but  only  a  single  kind  of  crystals 
which  are  entirely  different  in  nature  from  those  of  the 
two  simple  salts. 

The  present  preparation  illustrates  the  formation  of 
another  double  salt.  An-  equal  number  of  mols  of  copper 
sulphate  and  ammonium  sulphate  are  each  dissolved  in 
water,  and  on  mixing  the  solutions  there  are  produced 
crystals  of  the  formula  (NH4)2SO4.CuSO4.6H2O. 

Whether  or  not  a  double  salt  will  separate  from  a  mixed 
solution,  such  as  this  one,  depends  upon  the  solubility  rela- 
tions existing  among  the  various  compounds  capable  of 
formation.  In  this  case  the  double  salt  is  considerably 
less  soluble  than  either  of  the  simple  salts.  On  the  other 


7O  HEAVY    METALS 

hand,  if  saturated  solutions  of  sodium  sulphate  and  copper 
sulphate  are  mixed,  no  double  salt  will  separate,  but,  on 
evaporation,  a  mixture  of  the  crystals  of  the  two  simple 
salts  will  be  obtained.  Thus  it  appears  that  the  double 
sulphate  of  sodium  and  copper  is  more  soluble  than  the 
simple  salts. 

Procedure.  —  Weigh  out  33  grams  of  ammonium  sulphate 
and  62  grams  of  crystallized  copper  sulphate,  and  grind 
both  salts  together  in  a  mortar  until  they  are  very  finely 
powdered.  Add  the  mixture  to  250  cc.  of  water  in  a 
porcelain  dish ;  raise  the  temperature  to  just  5°  above  the 
room  temperature  and  keep  it  at  that  point  during  10  min- 
utes, while  stirring  all  of  the  time  with  the  stem  of  a  ther- 
mometer. This  gives  a  practically  saturated  solution  of  the 
double  salt.  Let  the  undissolved  salt  settle  a  moment  and 
pour  off  the  solution  into  another  dish.  With  a  little  more 
water  prepare  in  the  same  way  a  saturated  solution  of  any 
residue  that  is  left.  Add  a  few  drops  of  sulphuric  acid  to 
the  entire  solution,  warm  it  to  about  50°,  and  filter  it,  allow- 
ing the  filtrate  to  run  directly  into  an  8-inch  dish.  Smear 
a  thin  film  of  vaseline  around  the  inside  of  the  dish,  at  a 
little  distance  above  the  level  of  the  liquid,  in  order  to  pre- 
vent crusts  of  the  salt  from  creeping  up  over  the  edge  of  the 
dish.  Set  the  solution  away  uncovered  to  crystallize  in  a 
place  protected  from  the  dust ;  and,  before  it  has  cooled  to 
below  the  saturation  temperature,  seed  it  with  eight  or  ten 
small  crystals  of  the  double  salt.  Allow  the  solution  to 
stand  until  a  large  crop  of  crystals  has  formed.  Drain  the 
crystals,  and,  as  they  are  slightly  efflorescent,  leave  them 
spread  out  on  an  unglazed  dish  only  until  they  are  just  dry. 
If  only  a  small  amount  of  mother  liquor  is  left,  it  may  be 
discarded. 

It  is  a  matter  of  considerable  difficulty  to  obtain  clear, 
well-formed  crystals  of  this  salt,  especially  if  the  laboratory 


AMMONIO-COPPER    SULPHATE  /I 

temperature  varies  a  good  deal.  If  the  conditions  are  too 
unsatisfactory  it  is  better  to  give  up  the  attempt  to  get  dis- 
tinct crystals  and  to  prepare,  instead,  a  crystalline  meal: 
Dissolve  the  same  amounts  of  the  two  simple  salts  in  150  cc. 
of  hot  water ;  add  a  few  drops  of  sulphuric  acid ;  filter,  and 
cool  the  solution  rapidly  in  a  flask  while  rotating  it  under 
the  water  tap.  Drain  the  crystal  meal  on  a  Witt  filter  and 
dry  it  as  directed  above. 

Questions 

1.  Experiment.  —  Dissolve  a  little  of  the  salt  in  water. 
What  ions  does  the  solution  contain  ?     Make  the  ordinary 
tests  for  copper  and  sulphate  ions.      Does  the  double  salt 
exist  as  such  in  solution  ? 

2.  Compare  this  double  salt  with  the  double  cyanide 
of   silver   and   potassium   with   regard    to   its   tendency   to 
dissociate.     Experiment.  —  Make  a  little  of  the  latter  salt 
by  adding  to    i   cc.   of  silver  nitrate   solution   just  enough 
potassium   cyanide    solution    to    redissolve    the    precipitate 
which  at  first  forms.     Then  in  this  solution  of  potassium 
silver  cyanide  test  for  simple  silver  ions  by  adding  a  little 
sodium  chloride  solution  (not  hydrochloric  acid). 

3.  Find   out  by   experiment  whether   the  double    salt 
K2SO4.CuSO4.6H2O,  analogous  to  (NH4)2SO4.CuSO4.6H2O, 
can  be  prepared. 

15.     AMMONIO-COPPER    SULPHATE 

CuSO4.4NH3.H2O 

The  copper  and  ammonium  sulphate  of  the  last  prep- 
aration was  produced  by  the  combination  of  two  simple 
salts  and  is  therefore  a  so-called  double  salt.  Ammonio- 
copper  sulphate,  on  the  other  hand,  is  prepared  from  only 
one  simple  salt,  copper  sulphate,  which  in  crystallizing  is 


72  HEAVY    METALS 

caused  to  unite  with  four  molecules  of  free  ammonia  and 
one  molecule  of  water,  much  in  the  same  way  as  copper 
sulphate  unites  with  five  molecules  of  water  when  it  crys- 
tallizes as  blue  vitriol.  The  molecules  of  ammonia  would 
appear  to  be  bound  to  the  copper  atom  of  the  salt  rather 
than  to  the  sulphate  radical,  because  when  the  salt  is 
dissolved  in  water  the  four  ammonia  molecules  remain  in 
combination  with  the  copper,  giving  the  complex  ion 
Cu(NH3)4++,  while  the  sulphate  radical  gives  only  the 
ordinary  SO4~~-ion.  Thus  we  might  say  that  this  salt  is 
the  sulphate  of  the  ammonio-copper  complex. 

The  salt  is  exceedingly  soluble  in  water,  and  therefore, 
in  preparing  it,  use  is  made  of  its  insolubility  in  alcohol. 
The  method  adopted  of  allowing  the  alcohol  to  mix  with 
the  aqueous  solution  only  by  slow  diffusion  is  to  insure  the 
formation  of  large,  well-defined  crystals. 

Procedure.  —  Dissolve  25  grams  of  finely  pulverized  blue 
vitriol  in  84  cc.  of  dilute  ammonium  hydroxide  (sp.  gr.  0.96).1 
Place  125  cc.  of  alcohol  in  a  medium-sized  common  bottle. 
Fill  the  stem  of  a  small  separatory  funnel  with  water.  In- 
sert to  the  bottom  of  the  alcohol,  and  run  in  20  cc.  of 
water  to  form  a  layer  beneath  the  alcohol  to  separate  it 
from  the  ammoniacal  copper  solution,  which  is  next  intro- 
duced through  the  funnel.  Allow  no  bubbles  of  air  to  be 
sucked  with  the  liquid  into  the  stem  of  the  funnel  and  thus 
avoid  mixing  the  layers.  Set  the  bottle  away  for  at  least 
a  week,  at  the  end  of  which  time  crystals  2  and  3  cm.  long 
will  have  formed.  The  alcoholic  and  aqueous  layers  have 
not  yet  completely  diffused  into  each  other,  and  when  they 
are  mixed  a  meal  of  very  small  crystals  is  precipitated. 
Therefore  pour  the  liquid  all  at  once  out  of  the  bottle  into 
a  clean  beaker.  The  large  crystals  adhere  to  the  inside 
of  the  bottle.  Remove  them  to  a  small  dish;  add  10  cc.  of 

1  If  the  solution  is  not  perfectly  clear,  it  must  be  filtered  through 
asbestos  (Note  4  (</)),  since  it  dissolves  filter  paper  (cellulose). 


AMMONIO-COPPER    SULPHATE  73 

alcohol  to  which  i  cc.  of  ammonia  has  been  added ;  stir 
thoroughly  by  rotating  the  dish,  and  pour  off  the  alcohol, 
allowing  it  to  carry  with  it  any  of  the  precipitate  of  small 
crystals.  Repeat  the  washing  with  10  cc.  of  plain  alcohol 
and  then  with  10  cc.  of  ether.  Spread  the  crystals  on  filter 
paper  and  leave  them  until  they  cease  to  smell  of  ether. 
Then  put  them  up  at  once  and  stopper  them  tightly,  since 
they  give  off  their  ammonia  rather  easily.  Drain  on  a  Witt 
filter  the  crystal  meal  formed  in  the  beaker,  and  wash  it  on 
the  filter  (Note  5  (a))  with  the  same  liquids  as  were  used 
for  the  larger  crystals.  Preserve  the  large  crystals  and  the 
crystal  meal  separately. 

Questions 

1.  Experiments.  —  (a)  What   is   the   reaction   between 
Cu++   and    OH~-ions?      Add    sodium    hydroxide    solution, 
drop  by  drop,  to  a  solution  of  copper  sulphate  until  it  is 
present  in  large  excess.     Save  the  solution. 

(b~)  When  ammonium  hydroxide  is  added  instead  of 
sodium  hydroxide,  explain  what  successive  reactions  occur. 

(c)  Add  ammonium  chloride  to  the  solution  saved  from 
(a)  and  explain  the  results. 

(d)  To  5  cc.  of  copper  sulphate  solution  add  10  cc.  of 
a  10  per  cent,  solution  of  tartaric  acid;  then  add  sodium 
hydroxide  solution,  as  in  (a),  and  compare  the  results  with 
those  in  (a)  and  (£),  but  do  not  attempt  to  ascribe  a  definite 
formula  to  the  complex  compound  formed. 

2.  Experiments.  —  (a)  To   i   cc.  of  silver  nitrate  solu- 
tion add  sodium  hydroxide  until  present  in  large  excess. 

(b)  Repeat,  using  ammonium  hydroxide  instead  of  so- 
dium hydroxide,  and  adding  the  reagent  at  first  only  a 
single  drop  at  a  time.  Finally,  after  an  excess  of  5  cc.  of 
ammonium  hydroxide  is  present,  test  the  solution  with  a 
little  potassium  chloride  for  the  presence  of  silver  ions. 
Compare  the  results  with  those  in  Experiment  i. 


74  HEAVY    METALS 

1  6.     ZINC  OXIDE 

Zinc  oxide  is  used  as  a  white  pigment,  for  which  purpose 
it  has  the  advantage  of  not  turning  black  under  the  action 
of  hydrogen  sulphide.  It  can  be  obtained  directly  by  burn- 
ing metallic  zinc,  or  from  a  soluble  zinc  salt  by  precipitating 
first  a  basic  carbonate  and  then  heating  this  to  convert  it 
into  the  oxide.  Both  zinc  carbonate  and  zinc  hydroxide  are 
insoluble  in  water,  but  the  basic  carbonate  is  of  still  greater 
insolubility,  and  therefore  precipitates  more  readily  than  the 
former  two  when  the  ions  necessary  to  its  formation  are 
brought  together.  The  simplest  formula  for  the  basic 
carbonate  is 


but  the  precipitate  may  be  of  very  varying  composition 
according  to  the  conditions  of  its  formation. 

If  zinc  sulphate  in  solution  is  treated  with  sodium  bicar- 
bonate, pure  zinc  carbonate  is  precipitated,  because  a  sodium 
bicarbonate  solution  contains  but  a  minute  quantity  of  OH~- 
ions.  On  the  other  hand,  a  sodium  carbonate  solution,  in 
consequence  of  hydrolysis,  contains  a  considerable  quantity 
of  OH~-ions,  and  thus  it  furnishes  both  the  CO3~~  and  OH~- 
ions  necessary  for  the  formation  of  basic  zinc  carbonate. 

Zinc  carbonate  is  decomposed  by  heat  into  zinc  oxide 
and  carbon  dioxide. 

Procedure.  —  Dissolve  50  grams  of  crystallized  zinc  sul- 
phate (white  vitriol)  in  1.5  liters  of  hot  water  in  a  common 
2-liter  bottle,  and  add  slowly,  with  stirring,  a  solution  of 
19  grams  of  anhydrous  sodium  carbonate  in  250  cc.  of  hot 
water.  Let  the  precipitate  settle  somewhat,  and  test,  by 
adding  a  few  drops  more  of  sodium  carbonate  to  the  clear 
part  of  the  solution,  whether  all  the  zinc  has  been  precipi- 
tated. Let  settle  as  much  as  possible  (in  30  minutes  per- 
haps to  one-third  or  one-fourth  of  the  bulk  of  the  liquid). 


ZINC    OXIDE  75 

Draw  off  the  clear  liquid  and  wash  the  remaining  precipitate 
by  decantation  until  it  is  calculated  that  it  is  contaminated 
with  less  than  o.i  percent  of  the  soluble  sodium  sulphate 
present  at  first  (see  Note  5  (^)).  Finally,  transfer  the  slime 
to  a  large,  ordinary  filter  (Note  4  (r)),  and  allow  it  to  drain 
over  night.  Without  removing  the  pasty  product  from  the 
filter  open  out  the  latter  on  an  unglazed  plate,  and  leave  it 
on  the  steam  table  until  the  material  is  dry.  Pulverize  the 
basic  zinc  carbonate,  and  heat  it  in  a  small  porcelain  dish 
over  a  free  flame  until  all  the  carbon  dioxide  has  been 
driven  off  and  the  remaining  zinc  oxide  is  yellow  when  hot 
and  pure  white  when  cold. 

Questions 

1.  Why  could  not  the  precipitate  of  basic  zinc  carbon- 
ate have  been  advantageously  freed  from  the   solution  by 
means  of  a  suction  filter  ? 

2.  What  test   can    you   apply  to  prove  that   the  zinc 
carbonate    has    been    entirely    converted    into    the    oxide  ? 
Make  the  test. 

3.  Which  is  the  more  strongly  basic,  calcium  oxide  or 
zinc  oxide?     Which,  then,  could  be   more  readily  decom- 
posed by  heat,  calcium  carbonate  or  zinc  carbonate? 

4.  Experiment.  —  To    some   solution  of    zinc    sulphate 
add  a  solution  of  sodium  hydroxide,  drop  by  drop,  until  the 
precipitate  first  formed  redissolves.     How  is  zinc  hydroxide 
similar  to  aluminium  hydroxide   in  respect  to  its  behavior 
towards  strong  acids  and  strong  bases  ? 

5.  Experiment.  —  Knead  up  to  a  stiff   paste   about  a 
gram  of  zinc  oxide  with  a  few  drops  of  very  concentrated 
zinc  chloride   solution,  mold   into  a  lump,  and  observe  its 
condition    at   the    end    of    half    an    hour.      What   chemical 
change   occurs    when    this   mixture    hardens  ?     Compare    it 
with  the  change  which  occurs  when  plaster  of  Paris  "  sets," 


/6  HEAVY    METALS 

Add    a   few    drops   of    the   concentrated    zinc   chloride 
solution  to  a  tube  of  water.     What  is  the  precipitate  ? 


17.     MERCUROUS    NITRATE 
HgN03.H20 

Like  copper,  mercury  will  dissolve  in  nitric  acid,  an 
oxide  of  mercury  first  being  produced  in  consequence  of 
the  oxidizing  action  of  nitric  acid,  and  this  oxide  immedi- 
ately reacting  with  the  acid  to  form  a  salt.  Mercury  is 
capable  of  forming  two  oxides,  Hg2O  and  HgO,  and  two 
series  of  salts  derived  from  these  oxides  in  which  the  metal 
displays  the  valence  I  and  II,  respectively.  In  order  to 
obtain  the  salt  corresponding  to  the  lower  oxide,  it  is  only 
necessary  to  keep  an  excess  of  mercury  present,  through- 
out the  action  and  not  to  allow  the  nitric  acid  to  act  too 
violently. 

Procedure.  —  Treat  25  grams  of  mercury  in  a  flask  at 
the  hood  with  20  cc.  of  dilute  nitric  acid  (sp.  gr.  1.2).  warm- 
ing gently  until  no  further  action  takes  place.  Allow  to 
cool  until  the  flask  can  be  held  in  the  hand,  then  pour  the 
solution  away  from  any  remaining  globule  of  mercury  into 
a  small  dish,  and  leave  to  crystallize  until  the  next  day. 
Spread  the  crystals  out  upon  an  unglazed  porcelain  plate, 
and  stopper  them  in  a  test  tube  as  soon  as  dry. 

Questions 

i.  Test  the  preparation  for  mercuric  salt.  Treat 
0.5  gram  of  the  preparation  with  10  cc.  of  cold  water. 
It  will  not  dissolve  to  give  a  clear  solution.  Note  the 
character  of  the  residue.  Add  nitric  acid,  drop  by  drop, 
until  a  clear  solution  is  obtained.  By  adding  dilute  hydro- 
chloric acid,  drop  by  drop,  all  of  the  mercurous  salt  can 
be  precipitated,  leaving  any  mercuric  salt  in  solution,  Fil- 


MERCURIC    NITRATE  77 

ter ;  test  the  filtrate  for  mercuric  salt  by  means  of  hydrogen 
sulphide.  Judge  from  the  amount  of  the  precipitate  the 
quantity  of  mercuric  salt,  if  any. 

2.  Explain  why   mercurous    nitrate    does    not    dissolve 
completely  in  pure  water  and  what  is  the  character  of  the 
residue.       How    do    a    few    drops    of    nitric    acid    aid    in 
the  solution  ? 

3.  Explain  the  use  of  hydrochloric  acid  in  separating 
mercuric  from  mercurous  salt. 


18.     MERCURIC    NITRATE 

When  mercury  is  heated  with  an  excess  of  nitric  acid, 
mercuric  nitrate -is  produced.  This  salt  is  exceedingly  solu- 
ble in  water,  and  it  can  only  be  crystallized  with  a  good 
deal  of  difficulty.  When  a  solution  of  it  which  contains 
an  excess  of  nitric  acid  is  evaporated,  it  becomes  a  thick, 
heavy  sirup,  which  by  further  driving  off  of  nitric  acid  and 
water  becomes  a  pasty  mass,  due  to  formation  of  small 

NO 

crystals  of  basic  nitrate,  H£<CoH3'     If  the  materials  taken 

for  the  preparation  of  this  salt  are  pure,  the  product  can 
contain  no  other  foreign  matter  than  an  excess  of  nitric 
acid ;  consequently,  in  view  of  the  difficulty  of  obtaining 
good  crystals,  it  is  convenient  to  preserve  the  salt  in  this 
pasty  condition. 

Procedure.  —  Heat  25  grams  of  mercury  in  a  flask  with 
48  cc.  of  dilute  nitric  acid  (sp.  gr.  1.2)  until  it  is  all  dis- 
solved. Test  a  drop  of  the  solution  by  adding  to  it  in  a 
test  tube  \  cc.  of  cold  water  and  a  drop  of  dilute  hydro- 
chloric acid.  A  precipitate  will  probably  form,  in  which 
case  add  10  cc.  of  concentrated  nitric  acid  to  the  flask 
and  boil  until  a  precipitate  is  no  longer  obtained  when 
tested  as  above.  Pour  the  solution  into  a  dish  and  evapo- 


78  HEAVY    METALS 

rate  over  a  very  small  free  flame  until  the  liquid  has  assumed 
a  sirupy  consistence  and  crystals  just  commence  to  form  on 
the  surface.  Then  remove  the  whole  mass  to  a  small,  wide- 
mouth  sample  bottle,  which  has  previously  been  weighed ; 
let  it  cool  and  stopper  it. 

Questions 

1.  To  prepare  a  solution  of  this  salt  for  use  as  a  lab- 
oratory reagent,  explain  why  it  is  necessary  to  add  a  small 
quantity  of  nitric  acid.     (Compare  with  Questions  i  and  2 
under  Mercurous  Nitrate.) 

2.  Experiment.  —  To  a  solution  of  mercuric  nitrate  add 
a  little  hydrochloric  acid.     Now  add  a  liftle  stannous  chlo- 
ride solution.     What   is   the  precipitate,   and  what  change 
in  the  valence  of   mercury  must   have  occurred   before  it 
could  form  ? 


19.     MERCURIC    SULPHOCYANATE 

In  most  of  its  properties  the  sulphocyanate  radical 
resembles  the  halogens,  with  which  it  is  often  classed,  in 
the  same  manner  that  the  ammonium  radical,  NH4,  is 
classed  with  the  alkali  metals.  Thus  potassium  sulpho- 
cyanate, KSCN,  yields  the  ion  SCN~  just  as  potassium 
chloride  yields  the  ion  Cl~.  Mercuric  sulphocyanate  is 
insoluble  in  water,  and  can  be  produced  by  bringing  to- 
gether equivalent  quantities  of  solutions  of  mercuric  nitrate 
and  potassium  sulphocyanate,  but  if  an  excess  of  either  of 
these  reagents  is  used  the  precipitate  dissolves  in  it  with 
the  formation  of  a  complex  soluble  compound.  While  by 
a  consideration  of  the  Mass  Law  one  would  predict  that 
an  excess  of  either  ion  would  cause  a  decrease  in  the  solu- 
bility of  the  salt  Hg(SCN)2,  another  factor  nevertheless 


MERCURIC    SULPHOCYANATE  79 

comes  into  play,  namely,  the  tendency  of  the  molecule 
Hg(SCN)2  to  combine  either  with  Hg++  or  SCN~-ions  to 
form  complex  ions,  which  under  the  conditions  given  are 
capable  of  remaining  in  solution.  Similar  phenomena  are 
not  of  infrequent  occurrence  where  a  given  amount  of  a 
reagent  will  cause  a  precipitate,  while  an  excess  will  cause 
the  precipitate  to  redissolve.  For  instance,  carbon  dioxide 
gas,  if  passed  into  a  solution  of  calcium  hydroxide,  causes 
a  white  precipitate,  but  more  of  the  gas  causes  the  precipi- 
tate to  redissolve.  A  very  neat  expedient  may  be  adopted 
in  this  preparation  to  show  when  the  proper  amount  of 
reagent  has  been  added,  as  follows :  ferric  sulphocyanate, 
Fe(SCN)3,  is  a  soluble  substance  which  has  an  intensely  red 
color.  If  to  a  given  solution  of  mercuric  nitrate  a  few  drops 
of  a  ferric  salt  solution  are  added,  and  then  to  this  is  grad- 
ually added  a  solution  of  potassium  sulphocyanate,  the 
SCN~-ions  will  unite  with  the  Hg++-ions  so  long  as  any  of 
the  latter  are  present,  the  solution  remaining  colorless  and 
the  precipitate  Hg(SCN)2  forming  towards  the  end ;  but  so 
soon  as  the  Hg++-ions  are  exhausted,  then  the  SCN~-ions 
unite  with  Fe+  +  +-ions,  producing  the  red  compound,  and 
the  appearance  of  the  red  color  is  the  indication  to  stop. 

Mercuric  sulphocyanate  has  the  peculiar  property  that 
when  ignited  it  burns  with  the  production  of  a  very  volu- 
minous coherent  ash,  which,  from  the  form  which  it  assumes, 
is  called  "  Pharaoh's  Serpent."  It  should  not  be  burned 
in  an  open  room  because  of  the  production  of  mercury 
vapor,  which  is  poisonous. 

Procedure.  —  Dissolve  the  mercuric  nitrate  obtained  in 
the  last  preparation  in  i  liter  of  water,  adding  enough  nitric 
acid  to  prevent  the  formation  of  any  basic  salt.  To  this 
add  10  drops  of  a  ferric  chloride  solution;  then  add  grad- 
ually, with  constant  stirring,  a  solution  of  25  grams  of 
potassium  sulphocyanate  in  500  cc.  of  water  until  a  red 


8O  HEAVY    METALS 

color  appears  and  persists  after  stirring.  Collect  the  pre- 
cipitate on  a  Witt  filter,  and  dry  it  upon  an  unglazed  plate. 
The  dried  salt  may  be  made  into  the  so-called  Pharaoh's 
serpent  eggs  by  kneading  it  with  1.5  grams  of  dextrine 
dissolved  in  5  cc.  of  water  to  obtain  a  stiff  paste,  molding 
this  into  small  cones  or  pellets,  and  allowing  them  to  dry. 

Questions 

1.  Observations  made  in  the  course  of  this  prepara- 
tion are  sufficient  to  decide  whether  ferric  sulphocyanate, 
Fe(SCN)3,  is  an  ionized  or  un-ionized  substance.     Explain. 

2.  Find  out  what  is  the  degree  of  dissociation  of  the 
soluble  halides  of  mercury,  /.  <?.,  HgCl2,  Hg(CN)2.     Do  these 
salts  form  in  this  respect  any  exception  to  the  general  rule 
regarding  the  ionization  of  salts  ? 

3.  Describe  at  least  three  instances  which  have  pre- 
viously fallen  under  your   observation  in  which  a  reagent 
in  limited   amount  will   give  a  precipitate,  but,  added   in 
excess,  will  cause  the  precipitate  to  redissolve. 


GENERAL    QUESTIONS.     Ill 

HEAVY   METALS   OF  THE   FIRST  TWO    GROUPS   OF  THE   PERIODIC 
SYSTEM 

Experiments 

(The  results  observed  are  to  be  recorded  in  the  laboratory  note- 
book at  the  time  the  experiments  are  performed.) 

1.  Heat  a  little  dry  copper  carbonate  rather  gently  in 
a  test  tube  by  shaking  it  in  a  Bunsen  flame  until  the  color 
has  completely  changed  to  black.     Then  test  the  residue 
for  carbonate  by  adding  hydrochloric  acid. 

2.  To  solutions  of  copper,  silver,  zinc,  cadmium,  and 
mercury  salts,  in  separate  test  tubes,  add  sodium  hydroxide 


GENERAL    QUESTIONS    III  8 1 

until   a   precipitate   forms   in  each   case,  and   then  add   a 
considerable  excess  of  the  reagent. 

3.  Repeat   Experiment   2,  using  ammonium   hydroxide 
instead  of  sodium  hydroxide. 

4.  Place  small  pieces  of  copper  wire  in  solutions  of 
silver  and  mercury  salts.     Note  whether  any  metal  is  pre- 
cipitated on  the  surface  of  the  copper,  and  test  the  solutions 
by  adding  an  excess  of  ammonia  to  prove  whether  any  cop- 
per has  passed  into  solution.     Place  pieces  of   zinc  in  a 
copper  salt  solution,  let  it  stand  for  some  time  with  fre- 
quent shaking,  and  find  out  whether  all  of  the  copper  can 
be  precipitated  by  means  of  the  zinc. 

Questions 

1.  The  heavy  metals  of  the  first  two  groups  of  the  periodic  sys- 
tem show  certain  similarities   to   the  alkali  and  alkaline  earth  metals 
with  respect  to  the  type  of   compounds  formed.     In  this  way  their 
positions  in  the  same  groups  are  in  a  measure  justified.     They  show, 
however,  more   divergencies  in  their  properties,  and  on  this  account 
they  fall  naturally  into  separate  families  within  the  groups. 

Which  of  these  metals  are  capable  of  showing  different 
valences?  Give  the  symbols  of  the  chlorides  in  which 
the  valence  of  the  metal  is  characteristic  of  its  position 
in  the  group. 

2.  Compare    the    stability   of    copper    carbonate    (see 
Experiment   i)   with  that  of  potassium  carbonate ;  that  of 
zinc  carbonate  with   that  of   calcium  carbonate.     How  do 
the  heavy  metals  of  the  first  two  groups  compare  as  re- 
gards  basic    strength    with    the    alkali    and    alkaline    earth 
metals  ?     How    do    they    compare    with    aluminum    in    this 
respect  ? 

3.  What    in    general    terms   can   be    said    about    the 
solubility  of   the    hydroxides   of   copper,   silver,    zinc,  cad- 
mium,   and    mercury    (Experiment    2)  ?     Are    the    salts    of 
these  metals  much  or  little  hydrolyzed  in  aqueous  solution  ? 


82  HEAVY    METALS 

What    does   this    indicate    in    regard    to    the    base-forming 
characteristics  of  the  metals  ? 

4.  Which  of  the  hydroxides   of  these  metals  are  am- 
photeric?     (Experiment  2.) 

5.  Give  the  formulae  of  the  soluble  ammonio-salts  of 
copper,  silver,  and  zinc.     (Experiment  3.) 

6.  Arrange  the  metals  of  Family  B  in  Groups  I  and  II 
in  the  order  of  their  electrolytic   solution  tension  (Experi- 
ment 4).     Compare  Families  A  and  B  in  each  group  with 
respect  to  this  property. 


CHAPTER    IV 

ELEMENTS   OF   THE   FOURTH    GROUP   OF   THE 
PERIODIC   SYSTEM 

This  group  stands  in  the  middle  of  the  Periodic  Table 
of  the  elements,  and  in  it  the  difference  in  properties 
between  the  elements  of  Family  A  and  Family  B  is  at  a 
minimum.  As  in  the  case  of  Group  III,  therefore,  the 
whole  group  is  considered  under  one  heading. 

The  only  elements  of  this  group  which  are  to  be  espe- 
cially considered  are  carbon,  silicon,  tin,  and  lead.  Carbon 
and  silicon  are  the  first  two  members  and  are  exclusively 
acid-forming  elements,  although  the  acids  formed  are  not 
strong  ones.  Tin  and  lead  are  the  last  two  members  of 
Family  B  and  are  in  the  main  base-forming ;  they  are  com- 
parable in  this  respect  with  the  heavy  metals  already  consid- 
ered under  Groups  I  and  II.  It  is  a  characteristic  of  this 
group  that  the  acid-forming  properties  are  most  pronounced 
in  the  elements  of  low  atomic  weight,  while  the  base-forming 
properties  are  most  pronounced  in  the  elements  of  high 
atomic  weight. 

Titanium,  although  it  is  a  fairly  abundant  constituent 
of  the  earth's  crust,  is  an  element  of  comparatively  little 
importance,  and  is  never  taken  up  in  connection  with  the 
more  common  elements. 

Cerium  and  thorium,  the  heaviest  two  elements  of  Fam- 
ily A,  have  acquired  some  importance  on  account  of  the 
use  of  their  oxides  in  incandescent  gas-lighting  mantles. 


20.     STANNOUS    CHLORIDE    (SnCl2.2H2O) 

This  salt  can  be  prepared  by  the  action  of  hydrochloric 
acid  upon  metallic  tin.  Since,  however,  the  action  would 
be  exceedingly  slow,  it  may  be  hastened  by  the  addition 
of  a  very  small  quantity  of  nitric  acid,  which  will  oxidize 
the  tin.  Nitric  acid  is  ordinarily,  by  its  action  upon  a  metal, 
reduced  only  to  the  oxide  NO  ;  but  in  the  course  of  this 
preparation  no  red  fumes  of  oxides  of  nitrogen  will  be  found 
to  escape,  because,  under  the  influence  of  tin  and  stannous 
chloride,  the  reduction  does  not  stop  at  nitric  oxide,  but 
continues  to  the  lowest  possible  step,  which  is  ammonia  or 
in  this  case  its  salt,  ammonium  chloride.  Stannous  salts 
are,  under  the  influence  of  the  oxygen  of  the  air,  oxidized 
quite  readily  to  stannic ;  to  prevent  this  happening  during 
the  evaporation  of  the  solution,  an  excess  of  metallic  tin  is 
kept  in  the  liquid. 

Procedure.  —  Place  100  grams  of  feathered  tin  in  a  large 
casserole,  cover  with  175  cc.  of  concentrated  hydrochloric 
acid  (sp.  gr.  1.2),  and  add  (at  the  hood)  25  cc.  of -dilute 
nitric  acid  (sp.  gr.  1.2),  a  little  at  a  time,  during  a  period 
of  10  minutes.  Then  concentrate  the  solution,  by  boiling 
over  a  free  flame,  to  a  volume  of  90-100  cc.,  at  which 
point  a  crystal  scum  will  form  on  blowing  on  the  surface 
of  the  hot  liquid.  There  should  still  be  left  a  small  amount 
of  undissolved  metal.  If  at  any  time  during  the  evapora- 
tion all  the  tin  becomes  used  up,  add  a  little  more.  Prepare 
an  asbestos  filter  (Note  4  (//),  page  10),  moisten  it  with 
concentrated  hydrochloric  acid,  and  filter  the  concentrated 
stannous  chloride  solution  before  it  has  cooled  to  below 
60-70°.  Finally,  rinse  out  the  casserole  with  15  cc.  of  con- 
centrated hydrochloric  acid  and  pour  this  liquid  through 

85 


86        ELEMENTS  OF  THE  FOURTH  GROUP 

the  filter,  letting  it  mix  with  the  main  part  of  the  solution. 
If  during  the  filtration  the  liquid  stops  flowing,  due  to  crys- 
tals separating  in  the  filter,  add  5—10  cc.  of  boiling  water. 
Pour  the  solution  into  a  6-inch  evaporating  dish,  and  leave 
it  to  evaporate  slowly  at  room  temperature  in  a  place  ex- 
posed to  the  air  but  protected  from  dust.  (The  solubility 
of  stannous  chloride  decreases  very  rapidly  with  decreasing 
temperature.  Hence  it  is  advantageous  to  carry  out  the 
crystallization  in  as  cool  a  place  as  possible.)  When  a  good 
crop  of  crystals  has  formed,  pour  off  the  liquid  into  another 
dish,  spread  the  crystals  on  an  unglazed  dish,  and  allow 
them  to  dry.  It  is  to  be  remembered  that  stannous  chlo- 
ride is  extremely  soluble  in  water  and  that  the  composition 
of  the  mother  liquor  is  not  far  different  from  that  of  the 
crystals  of  SnCl2.2H2O  which  separate.  Heat  the  remain- 
ing solution  carefully  just  to  the  boiling  temperature,  but 
do  not  allow  it  to  boil  more  than  a  moment.  In  this 
way  sufficient  water  and  hydrochloric  acid  are  expelled  to 
allow  another  crop  of  crystals  to  form.1  Set  the  solution 
aside  to  cool  and  evaporate,  as  before,  and  collect  another 
crop  of  crystals.  By  repeating  this  process  once  or  twice 
more,  almost  the  entire  mother  liquor  should  be  used  up 
and  nearly  the  calculated  yield  of  stannous  chloride  should 
be  obtained. 

Questions 

i.  Explain  why  during  this  preparation  no  red  oxides 
of  nitrogen  are  seen  to  escape  in  consequence  of  the 
reduction  of  nitric  acid  by  the  metal. 

If  nitric  acid  is  reduced  to  NH3,  show  how  many  more 

1  If  too  much  hydrochloric  acid  is  expelled  by  the  evaporation  and 
an  indistinctly  crystalline  precipitate  of  basic  salt  separates  on  cooling, 
add  a  few  drops  of  hydrochloric  acid,  redissolve  the  salt  by  warming, 
and  set  the  solution  to  crystallize  again,  seeding  it  when  cold  with  a 
small  crystal  from  the  first  crop. 


STANNOUS    CHLORIDE  8/ 

equivalents  of  oxygen  it  will  yield  for  the  oxidation  of  the 
tin  than  if  it  were  reduced  only  to  NO. 

Experiment,  —  To  test  for  the  presence  of  ammonium 
salt  in  the  product,  take  about  i  gram  of  the  crystals ;  dis- 
solve in  10  cc.  of  water  in  a  small  beaker.  Add  sodium 
hydroxide  solution  until  the  precipitate  first  formed  redis- 
solves.  Place  over  the  beaker  a  watch  glass,  on  the  under 
side  of  which  is  stuck  a  piece  of  moistened  red  litmus 
paper.  Place  some  cold  water  in  the  hollow  of  the  watch 
glass,  and  warm  the  solution  in  the  beaker  very  gently. 
What  will,  if  it  occurs,  indicate  the  presence  of  ammonium 
salt,  and  why  ? 

2.  Experiment.  —  Dissolve  i  gram  of  stannous  chloride 
crystals  in  i  to  2  cc.  of  cold  water.     Then  add  a  consider- 
able amount  of  water.     What  is  the  precipitate  ?     What  can 
be  added  to  prevent  its  formation  ? 

3.  Experiment.  —  To  a  cold  solution  of  stannous  chlo- 
ride add  sodium  hydroxide  until  it  has  redissolved  the  pre- 
cipitate first  formed.     Write  the  reaction.     Save  the  solution. 

4.  Experiment.  —  Pour   the    solution    saved    from    Ex- 
periment 3  over  a  little  bismuth  hydroxide  on  a  filter  paper. 
(The  latter  can  be  precipitated  for  the  occasion.)     Compare 
the  action  with  that  of  stannous  chloride  on  mercuric  chlo- 
ride..   (See  Experiment  2  under  Mercuric  Nitrate.) 

5.  Experiment.  —  Prepare    a    very    concentrated    cold 
solution  of  sodium  stannite :  Dissolve   i  gram  of  stannous 
chloride  in  i  cc.  of  water.     Dissolve  a  small  lump  of  sodium 
hydroxide  in  its  own  weight  of  water,  and  add  this  solution, 
a  drop  at  a  time,  to  the  first  solution  —  cooling  all  the  while 
under  the  water  tap  —  until  the  precipitate  at  first  formed 
redissolves.     Then  heat  the  solution.     Compare  the  action 
with  that  in  Experiment  4. 

6.  In  preparing  a  solution  of  stannous  chloride  for  a 
laboratory  reagent,  what    is    the    necessity   of    adding   free 


88         ELEMENTS  OF  THE  FOURTH  GROUP 

hydrochloric   acid   and   of    placing   a  piece   of   metallic  tin 
in  the  bottle? 


21.     STANNIC    SULPHIDE    (MOSAIC    GOLD) 

Stannic  sulphide,  SnS2,  is  the  higher  sulphide  of  tin, 
and  can  be  prepared  by  direct  combination  of  the  metal 
or,  still  better,  of  the  lower  sulphide,  SnS,  with  sulphur. 
Under  ordinary  conditions  these  two  substances  will  not 
react  at  a  temperature  below  that  which  will  decompose 
stannic  sulphide.  If,  however,  they  are  mixed  with  am- 
monium chloride  the  presence  of  this  substance  makes 
possible  the  combination  at  a  lower  temperature,  and  also, 
since  it  absorbs  heat  by  its  volatilization,  it  prevents  the 
temperature  from  rising  too  high.  The  stannic  sulphide 
formed  in  this  way  appears  as  soft,  glistening,  yellow 
crystals.  It  is  used  as  a  bronzing  powder,  and  is  known 
under  the  name  of  mosaic  gold.  In  physical  properties  it 
is  very  different  from  the  stannic  sulphide  which  can  be 
precipitated  by  hydrogen  sulphide  from  a  solution  of  stannic 
chloride. 

Procedure.  —  Stannous  Sulphide.  Weigh  out  40  grams 
of  commercial  stannous  chloride,  SnCl2.2H2O,  into  a  casse- 
role, add  25  cc.  of  concentrated  hydrochloric  acid  (1.2), 
and  warm  the  mixture  gently  until  the  salt  is  dissolved. 
Dilute  the  solution  to  1,500  cc.  with  boiling  water  in  a 
2-liter  common  bottle,  and  pass  hydrogen  sulphide  (gen- 
erated in  a  small  generator  bottle  from  ferrous  sulphide, 
broken  into  small  lumps,  and  hydrochloric  acid ;  the  gas 
is  not  to  be  taken  from  the  laboratory  taps)  into  the  solu- 
tion until  all  of  the  tin  is  precipitated.  To  test  for  complete 
precipitation,  filter  off  a  small  sample  and  treat  it  with  more 
hydrogen  sulphide.  Wash  the  precipitate  twice  by  decanta- 
tion  (Note  5  (£),  on  page  13),  then  collect  it  on  a  filter 


STANNIC    SULPHIDE  89 

without  suction  (Note  4  (V)) ;  wash  it  twice  on  the  filter 
with  hot  water  from  the  wash  bottle,  and  allow  it  to  drain 
thoroughly.  Without  tearing  the  filter,  remove  it  carefully 
from  the  funnel  and  spread  it  out  on  an  unglazed  plate 
with  the  precipitate  still  on  it.  After  drying  it  thoroughly 
at  the  steam  table,  separate  the  brittle  lumps  of  stannous 
sulphide  from  the  paper  and  powder  them. 

Stannic  Sulphide.  —  Grind  together  thoroughly  20  grams 
of  the  stannous  sulphide  just  prepared,  10  grams  of  sulphur, 
and  8  grams  of  ammonium  chloride.  Bring  the  mixture 
into  a  small  flask  and  place  it  on  a  sand  bath,  forcing  it 
down  as  much  as  possible  into  the  sand.  Cover  the  mouth 
of  the  flask  with  an  inverted  test  tube  or  small  beaker,  and 
protect  the  sides  with  asbestos  paper.  Heat  the  sand  bath 
for  2  hours,  at  first  rather  moderately,  and  finally  so  that 
the  bottom  of  the  iron  pan  is  a  bright  red  in  the  center. 
Cool,  break  the  flask,  and  preserve  the  layer  of  stannic 
sulphide  which  is  found  at  the  bottom. 

Questions 

1.  Experiment.  —  Prepare  a  little  stannic  sulphide  by 
precipitation    with    hydrogen    sulphide   from    a    solution    of 
stannic  chloride.     Compare  this  with  the  stannic  sulphide 
prepared  above  as  regards  its  solubility  in  hydrochloric  acid. 

2.  Experiment.  —  In  each  of  two  test  tubes  place  about 
\  gram  of  powdered  stannous  sulphide,  and  fill  each  tube 
to  the  same  height,  about  half  full,  with  colorless  ammonium 
sulphide.     To  one  tube  add  a  few  small  lumps  of  roll  sul- 
phur.    Bring  the  contents  of  both  tubes  just  to  the  boiling 
temperature,  and  then  let  both  stand  with  occasional  shak- 
ing   and   further   gentle    heating   until   the  black  stannous 
sulphide  in  one  tube  has  dissolved.     Pour  off  about  2  cc. 
of  the  clear  liquid  from  this  tube,  dilute  it  with   15   cc.  of 
water,  and  just  acidify  the  solution  with  hydrochloric  acid. 


9O        ELEMENTS  OF  THE  FOURTH  GROUP 

Sodium  sulphostannite  is  so  unstable  a  salt  that  even 
if  it  could  be  formed  it  would  revert  at  once  into  sodium 
sulphide  and  stannous  sulphide.  Sodium  sulphostannate 
is  a  more  stable  salt. 

Explain  the  reaction  and  write  the  equation  for  the  for- 
mation of  the  sodium  sulphostannate.  Give  the  equation 
for  the  reaction  with  hydrochloric  acid. 

Explain  the  analogy  between  the  oxy-salt  sodium  stannate 
and  the  sulpho-salt  sodium  sulphostannate. 


22.     STANNIC    CHLORIDE    (ANHYDROUS) 

Anhydrous  stannic  chloride,  SnCl4,  is  a  mobile  and 
volatile  colorless  liquid  with  a  boiling  point  at  114°.  It 
fumes  very  strongly  in  the  air;  with  water  in  moderate 
amount  it  forms  crystalline  compounds,  but  with  more  water 
it  dissolves  to  a  clear  solution.  It  is  prepared  by  the  action 
of  dry  chlorine  gas  upon  metallic  tin. 

Apparatus.  —  This  preparation  is  only  to  be  attempted 
if  four  consecutive  hours  are  available  in  the  laboratory, 
and  even  then,  the  apparatus  should  be  assembled  as  much 
as  possible  at  a  previous  exercise. 

Arrange  a  chlorine  generator,  the  gas  from  which  is 
to  be  passed  through,  first  one  wash  bottle  containing  water, 
and  then  two  wash  bottles  containing  concentrated  sul- 
phuric acid.  Place  150  grams  of  tin  in  a  tubulated  retort, 
and  place  the  retort  on  a  sand  bath.  The  neck  of  the 
retort  should  pass  into  a  long  condenser,  and  the  latter 
should  empty  into  a  200  cc.  distilling  flask,  in  which  has 
been  placed  some  tin  foil.  Connect  the  side  arm  of  the 
flask  with  a  bottle  containing  sodium  hydroxide  solution 
to  absorb  the  waste  chlorine.  The  tube  entering  the  bot- 
tle should  not  dip  into  the  liquid,  but  should  reach  down 
to  near  its  surface ;  a  safety  tube  should  be  supplied,  and 


STANNIC    CHLORIDE  9 1 

the  exit  tube  should  dip  into  a  solution  of  sodium  hydroxide 
in  another  bottle0  Into  the  retort  should  be  fitted  the  tube 
supplying  chlorine  from  the  generator  and  wash  bottles, 
and  this  should  reach  nearly  to  the  center  of  the  surface 
of  the  tin,  which  is  to  be  melted  before  the  action  is  started. 
Glass  tubing  is  to  be  used  throughout,  and  where  connec- 
tions are  made  with  rubber  the  ends  of  the  glass  tubes 
should  be  brought  close  together^  Before  beginning  to 
generate  the  chlorine,  the  whole  apparatus  must  be  proved 
to  be  tight,  so  that  none  of  this  gas  can  escape  into  the 
laboratory. 

L  Procedure.  —  Melt  the  tin.  Commence  the  generation 
of  chlorine  and  regulate  it  so  that  the  tin  in  the  retort  can 
be  seen  to  burn  quietly,  but  do  not  allow  this  action  to 
become  too  violent.  Continue  the  action  until  all  the  tin 
has  disappeared  and  the  tin  tetrachloride  has  been  caught 
in  the  receiving  flask.  Remove  the  neck  of  the  retort 
from  the  condenser,  and  insert  instead  a  stopper  with  a 
tube  leading  to  the  bottles  already  used  for  absorbing  waste 
chlorine.  Close  the  side  arm  of  the  receiving  flask  and, 
with  the  condenser  still  in  the  same  position,  boil  the  tin 
tetrachloride  until  it  is  colorless  (it  contains  a  large  amount 
of  dissolved  chlorine  and,  on  boiling,  this  reacts  with  the 
tin  foil).  Change  the  position  of  the  flask  and  condenser, 
and  distill  the  tin  tetrachloride  into  a  flask  whose  neck 
has  been  drawn  down  so  that  it  can  be  readily  sealed 
off.  During  the  distillation  the  flask  should  not  be  open 
to  the  air,  but  should  be  connected  by  a  tube  to  the 
absorbing  bottles  already  used.  When  the  liquid  is  all 
distilled,  seal  the  neck  of  the  flask  at  a  blast  lamp,  so  that 
the  preparation  can  be  preserved  out  of  contact  with  the 
air. 

Questions 

i.     What  is  the  purpose  of  the  wash  bottles  as  arranged 
for  the  chlorine  gas  ? 


92 


ELEMENTS  OF  THE  FOURTH  GROUP 


2.  Compare  the  physical  properties  of  the  tetrachlorides 
of  carbon,  silicon,  tin,  and  lead. 

How  do  these  substances  behave  in  presence  of  water, 
and  what  information  does  this  give  as  to  the  metallic 
character  of  the  four  elements  ? 


23.     LEAD,  NITRATE 

Lead  nitrate  is  one  of  the  most  readily  prepared  salts 
of  lead,  since  it  is  of  moderate  solubility  and  can  be  ob- 
tained in  well-formed  anhydrous  crystals  (Pb(NO3)2).  In 
it  lead  appears  in  its  usual  state  of  oxidation,  which  cor- 
responds to  that  of  the  oxide  PbO ;  indeed,  the  salt  is 
actually  prepared  by  treating  this  oxide  (litharge)  with  nitric 
acid. 


A  saturated  solution  contains  for  each  100  grams  of  water  the  given 
number  of  grams  of  lead  nitrate 


0° 

10° 

18° 

25° 

50° 

100° 

36 

44 

51 

56 

79 

127 

Procedure.  —  Take  50  grams  of  litharge.  Calculate  the 
amount  of  nitric  acid  which  would  be  necessary  to  convert 
it  into  lead  nitrate  and  the  amount  of  water  needed  to  dis- 
solve the  salt  which  is  thus  formed.  Proceed  to  prepare 
lead  nitrate,  striving  to  obtain  good  crystals  of  as  large  a 
size  as  possible. 

Questions 

1.  Explain  why  lead  nitrate  should  be  less  soluble  in 
dilute  nitric  acid  than  in  pure  water. 

2.  Add  a  few  drops  of  ammonium  hydroxide  to  i  cc. 
of  lead  nitrate  solution.     Then  add  an  excess  of  the  reagent 
Repeat,  using  sodium  hydroxide  instead  of  ammonium  hy- 


LEAD    DIOXIDE  93 

droxide.     Give  equations  and  explain  the  amphoteric  char- 
acter of  lead  hydroxide. 

3.  Are  lead  salts  (nitrate  or  chloride)  appreciably  hydro- 
lyzed  in  aqueous  solution  ?     Compare  the  basic  strength  of 
lead  hydroxide  with  that  of  aluminum  hydroxide. 

4.  Precipitate  a  little  lead   chloride  by  adding  hydro- 
chloric  acid   to   a  solution  of   lead   nitrate.      Describe    its 
properties  and  compare  them  with  those  of  lead  tetrachlo- 
ride  (text-book).     To  what  oxide  of  lead  does  lead  tetra- 
chloride  correspond  ? 


24.     LEAD    DIOXIDE 

As  already  stated,  the  most  frequently  occurring  and 
the  most  stable  compounds  of  lead  are  those  derived  from 
the  oxide  PbO.  Under  certain  conditions,  however,  this 
oxide,  or  any  salt  derived  from  it,  may  be  converted  into 
the  higher  oxide  PbO2.  Thus  if  a  solution  of  lead  acetate 
is  made  alkaline  and  is  then  treated  with  an  oxidizing  agent, 
such  as  chlorine,  the  dioxide  results.  On  account  of  the 
disagreeableness  of  using  chlorine  gas  in  the  laboratory, 
bleaching  powder,  which  exerts  practically  the  same  oxidiz- 
ing action,  will  be  substituted  for  it.  The  precipitate  finally 
obtained  after  the  bleaching  powder  has  acted  contains  the 
greater  part  of  the  lead  in  the  form  of  lead  dioxide ;  but  it 
may  also  contain  a  small  residue  of  unoxidized  lead,  as 
Pb(OH)2,  as  well  as  calcium  hydroxide  from  the  bleaching 
powder  and  possibly  the  salts  calcium  plumbate  and  calcium 
plumbite,  CaPbO3  and  CaPbO2.  By  treating  this  precipi- 
tate with  nitric  acid  everything  except  the  lead  dioxide  is 
dissolved  or  decomposed,  and  practically  pure  lead  dioxide 
remains. 

Procedure.  —  Dissolve  100  grams  of  lead  acetate  in  200  cc. 
of  cold  water  in  a  750  cc.  casserole;  add  a  solution  of  21 


94        ELEMENTS  OF  THE  FOURTH  GROUP 

grams  of  sodium  hydroxide  dissolved  in  100  cc.  of  water, 
stirring  well,  and  into  the  mixture,  which  should  not  be 
warmer  than  30°,  stir  a  paste  made  by  rubbing  60  grams 
of  bleaching  powder  in  a  mortar  with  a  little  water.  Warm 
the  mixture  slowly  to  the  boiling  temperature,  stirring  fre- 
quently, and  finally  boil  it  for  10  minutes.  Transfer  the 
contents  of  the  casserole  to  a  2-liter  common  bottle  and 
wash  the  precipitate  by  decantation  with  cold  water  (see 
Note  5  (b)  on  page  13)  until  the  wash  water  gives  no  test 
for  chlorine  ions.  Then  transfer  the  remaining  slime  again 
to  a  casserole,  add  250  cc.  of  nitric  acid  (1.2),  and  boil  it 
for  10  minutes.  Wash  the  residue  of  lead  dioxide  by  decan- 
tation until  the  wash  water  is  no  longer  acid ;  transfer  the 
product  to  a  filter  and  let  it  drain  without  suction  (Note  4 
(<r)  on  page  9).  After  the  lead  dioxide  has  drained,  remove 
the  filter  and  contents  carefully  from  the  funnel,  unfold  the 
filter,  and  spread  it  on  an  unglazed  plate  on  the  steam  table 
to  dry.  When  completely  dry,  detach  the  lumps  of  lead 
dioxide  from  the  paper,  and  pulverize  them  in  a  mortar. 

Questions 

1.  Experiment.  —  Add  sodium  hydroxide  to  a  solution 
of  a  lead  salt,  and  observe  with  what  readiness  the  precipi- 
tate redissolves  in  an  excess  of  the  reagent. 

Give  reactions,  and  compare  Pb(OH)2  with  A1(OH)8  as 
regards  the  strength  of  their  acidic  properties. 

2.  Write   the  reaction   between   sodium    plumb ite  and 
chlorine  in  the  production  of  lead  dioxide ;  between  sodium 
plumbite  and  bleaching  powder. 

3.  Show  the  relation  among  ortho-plumbic  acid,  meta- 
plumbic  acid,  and  lead  dioxide.     Give  the  symbol  of  sodium 
meta-plumbate  ;  of  calcium  ortho-plumbate. 

4.  In  view  of  the  following  experiment,  why  could  not 
lead  dioxide  have  been  prepared  equally  well  by  treating 


RED    LEAD  95 

a  solution  of  lead  chloride,  containing  also  free  hydrochloric 
acid,  with  chlorine?  Experiment.  —  Drop  a  grain  of  lead 
dioxide  the  size  of  a  pinhead  into  \  cc.  of  hydrochloric  acid, 
and  warm  until  dissolved.  What  odor  is  observed  ?  Cool 
the  liquid  and  observe  the  character  of  the  precipitate. 

5.  Compare  the  reaction  of   lead  dioxide  and  of   lead 
monoxide  upon  hydrochloric  acid. 

6.  Compare   the    action    of   lead    dioxide   upon   hydro- 
chloric acid  with  that  of  manganese  dioxide. 

7.  Why  should  not  lead  dioxide  and  manganese  dioxide 
dissolve  in  nitric  acid  as  well  as  in  hydrochloric  acid  ? 

25.     RED    LEAD,    Pb3O4 

Of  the  oxides  of  lead,  the  monoxide  PbO  is  the  most 
stable  when  heated  to  a  high  temperature,  and  in  fact  all 
of  the  other  oxides  are  converted  into  this  one  when  they 
are  heated  strongly  in  contact  with  the  air.  At  a  moderate 
heat,  however,  the  monoxide  is  capable  of  taking  on  more 
oxygen  from  the  air  until  the  composition  approximates 
that  of  the  formula  Pb3O4.  This  substance  is  not  to  be 
regarded  as  a  simple  oxide  of  lead,  but  rather  as  a  com- 
pound of  PbO  and  PbO2,  in  which  the  monoxide  is  the 
basic  component  and  the  dioxide  the  acidic.  It  may  thus 
be  regarded  as  the  salt,  lead  orthoplumbate,  2PbO.PbO2  = 
Pb2(PbO4).  This  view  is  strengthened  by  the  behavior  of 
the  substance  when  treated  with  nitric  acid  —  part  of  the 
lead  dissolves  to  give  lead  nitrate,  while  the  other  part  is 
left  as  lead  dioxide, 


Pb2(Pb04)  +  4HN08-*  2Pb(N03)2  +  H4(Pb04) 
H4PbO4  -  PbO2  -f  2H2O. 

The  following  procedure  should  yield  a  product  of  nearly 
the  composition  Pb3O4.  This  substance,  under  the  commer- 
cial name  of  minium,  finds  use  as  a  red  pigment. 


g         ELEMENTS  OF  THE  FOURTH  GROUP 

Procedure. — Spread  25  grams  of  lead  monoxide  in  a 
thin  layer  on  an  iron  or  aluminum  plate  2-4  mm.  thick. 
Either  use  the  variety  of  lead  oxide  which  has  not  been 
fused  and  is  known  under  the  name  of  massicot,  or  use  lead 
carbonate,  which  on  being  heated  yields  a  very  pure  and 
finely  divided  lead  monoxide.  Heat  the  lead  oxide  over  a 
ring  burner  so  adjusted  that  the  flames  do  not  quite  touch 
the  metal  plate.  The  latter  must  be  kept  just  below  a  per- 
ceptible red  heat.  Continue  the  heating  for  2  hours  or 
more  and  turn  over  the  powder  frequently  with  an  iron 
spatula.  When  the  change  is  complete,  the  product  is  dark 
brown  when  hot,  a  bright  scarlet-red  when  partly  cooled, 
and  a  somewhat  less  brilliant  red  when  entirely  cold. 

Questions 

1.  If   it  is   assumed  that  Pb2O3   and   Pb3O4  are    lead 
metaplumbate  and    lead   orthoplumbate,   respectively,   write 
formulae  to  express  these  facts.     Write  the  formulae  of  the 
corresponding  meta-  and  ortho-plumbic  acids. 

2.  Experiment.  —  Boil  a  little  of  the  red  lead  with  nitric 
acid.     What  is  the  residue,  and  what  soluble  salt  is  formed  ? 
Test  the  solution   by  diluting  and   adding  a  few  drops  of 
sulphuric  acid. 

3.  Experiment.  —  Heat  a  little  of  the  red  lead  to  a  dull 
red  heat  on  a  thin  piece  of  iron. 

GENERAL    QUESTIONS.     IV 

ELEMENTS   OF   THE   FOURTH    GROUP    OF   THE   PERIODIC   SYSTEM 

Experiments 

(The  results  observed  are  to  be  recorded  in  the  laboratory  note- 
book at  the  time  the  experiments  are  performed.) 

i.  Heat  a  little  lead  carbonate  gently  by  shaking  it  in 
a  test  tube  in  a  Bunsen  flame  until  its  color  has  changed 


GENERAL    QUESTIONS    IV  97 

entirely  from  white  to  brownish  yellow.     After  cooling,  test 
the  residue  for  carbonic  acid  by  adding  dilute  nitric  acid. 

2.  Heat  a  little  piece  of  tin  in  a  casserole  with  some 
nitric  acid   until  the   metal  has   all  been  converted  into  a 
white  powder.     What  is  the  powder  ? 

3.  Boil  a  little  litharge,  PbO,  in  a  test  tube  of  water 
for  a  few  minutes.     Then  test  the  solution  with  red  litmus. 
Also   filter  off  some   of  the   clear   solution  and   test   it  for 
lead  by  adding  hydrogen  sulphide  water. 

4.  To  a  little   diluted   stannous   chloride   solution  add 
sodium   hydroxide    solution,   drop  by  drop,   until    a    heavy 
white  precipitate   has  formed,  then  add   it  in  considerable 
excess.     Repeat,  using  a  solution  of  a  lead  salt  instead  of 
stannous    chloride,    and    observe    how   great   an    excess    of 
sodium  hydroxide  is  necessary  to  redissolve  the  precipitate 
at  first  formed. 

Questions 

1.  Of  the  elements  of  the  fourth  group  only  four  need 
be  given  especial  consideration,  namely,  carbon,  silicon,  tin, 
and   lead.     What  is  the  characteristic  valence   of  the   ele- 
ments of  this  group  ?     Give  the  symbols  of  all  the  oxides 
of  each  of  these  elements,  and  indicate  in  each  case  whether 
the  substance  is  gaseous  liquid  or  solid.     In  the  same  way 
give  the  symbols   and  indicate  the  state  of  aggregation  of 
the  chlorides  of  these  elements  when  in  the  pure  state. 

2.  Is  tin  carbonate  capable  of  existence  ?     What  is  the 
stability  of  lead  carbonate?     (See  Experiment  i.)     Are  salts 
of  carbon  and  silicon,  such  as  sulphates  or  nitrates,  capable 
of  existence  ?     How  does  the  strength  of  the  base-forming 
tendency  change  in  this  group  in  going  from  carbon  to  lead  ? 

3.  Is  a  higher  carbonate  of  lead  than  PbCO3  capable 
of    existence  ?      Can    lead   tetranitrate,    Pb(NO3)4,    be    pre- 
pared ?     How  does  lead  tetrachloride  behave  when  treated 
with  a  large  amount  of  water  ?     Which  of  the  oxides  of  tin 


98         ELEMENTS  OF  THE  FOURTH  GROUP 

is  capable  of  forming  the  corresponding  nitrate  and  sul- 
phate ?  (See  Experiment  2.)  Explain  how  these  facts 
show  whether  the  monoxides  or  the  dioxides  of  lead  and 
tin  are  the  more  strongly  base-forming. 

4.  Discuss  the  solubility  and  the  degree  of  the  electro- 
lytic dissociation  of  lead   hydroxide  (Experiment  3).     Dis- 
cuss  the  amphoteric  character  of    Sn(OH)2  and  Pb(OH)2 
(Experiment  4). 

5.  How    may    sodium    sulphostannate    be    prepared? 
What  relation  does  it  bear  to  sodium  stannate  ? 

6.  Of  the  five  distinct  oxides  of  lead  described  in  the 
text-books,  which  are  the  salt  formers,  /.  e.,  which  are  the 
simple  basic  or  acidic  anhydrides  ? 


CHAPTER    V 

ELEMENTS   OF   THE   FIFTH    GROUP   OF   THE 
PERIODIC    SYSTEM 

In  this  group,  as  in  Groups  III  and  IV,  the  difference 
in  properties  between  the  elements  of  Families  A  and  B  is 
not  so  striking  as  in  Groups  I  and  II  (or  as  in  Groups  VI 
and  VII).  Consequently  the  whole  group  is  considered 
under  the  same  heading.  But  it  is  also  true  that  the  ele- 
ments of  Family  A,  that  is,  vanadium,  columbium,  and  tan- 
talum, are  of  comparatively  infrequent  occurrence,  and  are 
given  no  attention  in  this  course.  On  the  other  hand  all 
of  the  elements  of  Family  B  are  of  frequent  occurrence  and 
considerable  importance. 

The  characteristic  valence  of  the  group  is  five,  corre- 
sponding to  the  oxide  M2O5,  but  the  elements  likewise  all 
exhibit  a  valence  of  three  in  the  oxide  M2O3.  It  is  note- 
worthy that  the  valence  is  nearly  always  either  three  or  five. 

It  is  true  in  this  group,  as  well  as  in  Group  IV,  that  the 
acid-forming  properties  are  most  marked  in  the  elements 
of  low  atomic  weight  (nitric  acid  is  one  of  the  strongest 
acids),  and  decrease  with  increasing  atomic  weight ;  whereas 
the  base-forming  properties  are  most  strongly  developed 
with  the  elements  of  high  atomic  weight. 


99 


26.     ORTHO-PHOSPHORIC   ACID.     H3PO4 

A  rather  impure  grade  of  phosphoric  acid  can  be  obtained 
from  natural  calcium  phosphate  by  decomposition  with  sul- 
phuric acid,  but  a  pure  product  may  be  most  readily  obtained 
by  oxidizing  phosphorus  by  means  of  nitric  acid. 

The  solution  first  obtained  by  the  action  of  dilute  nitric 
acid  upon  phosphorus  contains  a  considerable  quantity  of 
phosphorous  acid,  H3PO3 ;  but  upon  boiling  down  this  solu- 
tion, a  point  is  reached  at  which  a  rather  strong  reaction 
takes  place,  which  consists  of  an  oxidation  of  the  phos- 
phorous to  phosphoric  acid  by  means  of  the  nitric  acid 
still  present. 

Commercial  phosphorus  often  contains  a  small  quantity 
of  arsenic.  This  on  the  treatment  with  nitric  acid  is  oxidized 
to  arsenic  acid,  which,  unless  removed  by  hydrogen  sulphide, 
would  contaminate  the  preparation  of  phosphoric  acid. 

Phosphorous  acid  may  always  be  present  in  the  product 
in  case  the  oxidation  with  nitric  acid  has  not  been  complete, 
and  its  presence  may  be  detected  by  its  ability  to  reduce 
silver  nitrate  and  give  a  black  precipitate  of  metallic  silver, 

H3PO3  +  2AgNO3  +  H2O  =  2HNO3  +  2Ag  +  H3PO4. 

Procedure.  —  Place  275  cc.  of  nitric  acid  of  sp.  gr.  1.2 
in  a  2-liter,  round-bottom  flask.  Weigh  out  30  grams  of 
red  phosphorus,  and  at  the  hood  add  10  grams  to  the  nitric 
acid  in  the  flask.  Warm  by  placing  the  flask  in  a  pail  of 
hot  water  just  until  red  vapors  begin  to  appear ;  then  stand 
the  flask  on  a  wooden  ring,  and  place  a  large  funnel  in  its 
neck  to  condense  and  allow  to  drip  back  at  least  a  part 
of  the  vapors.  After  the  action  has  nearly  ceased,  add  a 
little  more  of  the  phosphorus  and  again  wait  until  action 
has  almost  stopped,  and  so  on  until  all  of  the  phosphorus  is 
used.  When  the  action  following  the  last  addition  of  phos- 

101 


1:02  E^IMIZNTS    OF    THE    FIFTH    GROUP 

phorus  has  somewhat  moderated,  add  20  cc.  of  concentrated 
nitric  acid  (sp.  gr.  1.42).  Boil  the  solution  in  a  casserole 
until  it  is  considerably  concentrated  and  a  rather  violent 
reaction  is  observed  to  take  place  with  evolution  of  red 
fumes.  When  this  action  has  ceased,  the  remaining  solu- 
tion is  poured  into  a  flask,  diluted  to  about  a  liter  with 
water,  saturated  with  hydrogen  sulphide  gas,  stoppered, 
and  allowed  to  stand  over  night.  If  the  next  morning  the 
contents  of  the  flask  smell  strongly  of  hydrogen  sulphide, 
the  precipitate  of  arsenic  sulphide  is  filtered  off ;  if  not,  the 
solution  is  again  treated  with  hydrogen  sulphide  in  the  same 
manner  as  before.  Evaporate  the  filtrate  until  its  tempera- 
ture has  risen  to  125°.  Test  a  few  drops  for  phosphorous 
acid  by  diluting  with  a  little  water,  adding  some  silver 
nitrate,  and  warming.  If  any  is  found  add  20  cc.  of  nitric 
acid  to  the  solution  and  warm  it.  Transfer  the  solution 
to  a  small  casserole  and  evaporate  it  over  a  very  small  flame 
until  a  thermometer  whose  bulb  is  immersed  in  it  stands 
at  1 80°.  During  this  final  evaporation  one  must  give  it 
constant  attention,  for  if  it  is  left  and  the  temperature  rises 
above  180°,  not  only  does  the  ortho-phosphoric  acid  become 
changed  partially  into  pyrophosphoric  acid,  but  it  attacks 
very  strongly  the  material  of  the  dish,  and  the  preparation 
becomes  contaminated  with  silicic  acid.  Pour  the  liquid 
while  still  warm  into  a  previously  weighed  small  glass- 
stoppered  bottle,  and  stopper  it  tightly.  When  cool  intro- 
duce a  small  crystal  of  phosphoric  acid  to  induce  crystalli- 
zation of  the  mass. 

NOTE.  —  If  no  crystallized  phosphoric  acid  is  obtainable  the  sirupy 
acid  obtained  above  can  be  made  to  crystallize  spontaneously  if  it  is 
placed  in  a  vacuum  desiccator  over  concentrated  sulphuric  acid  and 
cooled  with  a  freezing  mixture. 

Questions 

i .  Write  the  reaction  by  which  phosphoric  acid  can  be 
prepared  from  calcium  phosphate, 


ARSENIC    ACID  1 03 

2.  How  can  phosphoric  anhydride,  pyrophosphoric  acidt 
and    metaphosphoric    acid    be    prepared?       Give    symbols. 
Why  cannot  the  anhydride  be  prepared  by  heating  ortho- 
phosphoric  acid  ?     For  what  practical  purpose  is  phosphoric 
anhydride  used? 

3.  .  Compare  the  acid  strength  of  phosphoric  acid  with 
that  of  other  common  acids.     Do  all  three   hydrogen  ions 
dissociate  with  equal  readiness  ? 

4.  Give  the  symbols  of  primary,  secondary,  and  tertiary 
sodium  phosphates.     State  how  the  solution  of  each  behaves 
with  litmus. 

5.  Write  the  reaction  for  the  precipitation  which  occurs 
when  magnesium  chloride  is  added  to  a  solution  of  phos- 
phoric acid  or  a  phosphate,  which  is  made  strongly  ammoni- 
acal.     This  precipitate  constitutes  one  of  the  most  important 
tests  for  phosphoric  acid. 

6.  Give   an   example   of   how   phosphorous   acid   may 
behave  as  a  reducing  agent. 


27.     ARSENIC    ACID    (H3AsO4)2.H2O 

Arsenic  acid  in  its  properties  shows  a  striking  similarity 
to  phosphoric  acid ;  and  even  the  method  of  its  preparation 
is  similar,  in  that  use  is  made  of  the  oxidizing  action  of 
nitric  acid.  Instead  of  starting  with  uncombined  arsenic, 
however,  use  is  made  of  arsenious  oxide,  As2O3,  a  product 
which  condenses  in  the  flues  wherever  ores  which  contain 
arsenic  are  roasted.  By  the  nitric  acid  this  is  oxidized  to 
the  higher  oxide,  As2O5,  which,  with  water,  yields  arsenic 
acid,  H3AsO4.  From  the  salts  which  arsenic  acid  forms,  it 
is  known  to  be,  like  phosphoric  acid,  a  tribasic  acid,  that  is, 
one  which  yields  three  hydrogen  ions.  By  evaporating  its 
solution  for  a  long  time  on  the  water  bath,  crystals  of  ortho- 
arsenic  acid  having  the  composition  H8AsO4  can  be  oth 


IO4  ELEMENTS    OF    THE    FIFTH    GROUP 

tained.  By  prolonged  evaporation  at  higher  temperatures 
crystals  of  the  composition  H4As2O7  and  HAsO3,  respec- 
tively, can  be  obtained.  When  a  solution  of  arsenic  acid 
is  boiled  down  according  to  the  following  directions,  a  liquid 
is  obtained  of  almost  exactly  the  composition  given  by  the 
formula  (H3AsO4)2.H2O.  This  liquid  when  cooled  to  35.5° 
or  below  can  be  crystallized  to  a  solid  product  of  the  same 
composition,  and  this  is  the  most  satisfactory  form  in  which 
to  crystallize  arsenic  acid.  It  is  interesting  to  note  that  this 
liquid  can  be  much  supercooled  below  35.5°,  but  that  when 
once  crystallization  is  induced  the  temperature  immediately 
rises  to  this  point  and  remains  there  until  solidification  is 
complete.  Likewise  when  the  solid  is  being  melted  the 
temperature  will  not  rise  above  the  melting  point,  35.5°, 
until  the  whole  mass  is  liquefied. 

Procedure.  —  Place  50  grams  of  arsenious  oxide  in  a 
good-sized  casserole;  add  20  cc.  of  water,  and  then  at  the 
hood  add  75  cc.  of  concentrated  nitric  acid;  warm  occasion- 
ally just  enough  to  keep  up  the  action,  but  do  not  allow  the 
reaction  to  become  violent,  because  the  heat  would  drive 
off  the  nitric  acid.  Finally,  after  red  fumes  cease  to  be 
evolved,  evaporate  the  liquid,  holding  the  casserole  in  the 
hand  and  rotating  it  all  the  while,  until  the  residue  is  just 
dry.  When  cool  enough  add  60  cc.  of  water,  and  heat  until 
the  residue  is  completely  dissolved.1  Evaporate  the  solution 
by  boiling  it  gently  in  a  small  casserole  until  the  tempera- 
ture has  risen  to  115°.  Then  transfer  the  liquid  to  a  very 
narrow  beaker  or  test  tube,  and  boil  it  carefully  with  a  small 
flame  until  the  temperature  shown  by  a  thermometer  inserted 

1  If  complete  solution  does  not  take  place  after  heating  for  some 
minutes,  either  the  oxidation  of  the  arsenious  oxide  has  not  been 
complete  or  else  the  residue  has  been  heated  too  strongly.  Test  for 
arsenious  acid  according  to  the  directions  in  Experiment  5.  If  any  is 
found  it  must  be  oxidized  by  the  addition  of  more  nrtric  acid,  and  the 
liquid  must  again  be  evaporated  to  dryness. 


ARSENIC    ACID  10$ 

in  the  liquid  has  just  risen  to  160°.  Cool  the  product  to 
below  35.5°,  place  it  in  a  weighed  sample  bottle,  and  seed 
it  with  a  small  crystal  of  (H3AsO4)2.H2O,  whereupon  the 
whole  will  slowly  crystallize  to  a  solid  mass.  Stopper  the 
bottle  tightly,  since  arsenic  acid  takes  moisture  rapidly  from 
the  atmosphere. 

Questions 

1.  Give  the  reaction  of  nitric  acid  upon  arsenious  oxide. 

2.  Compare  the  strength  of  arsenic  acid  with  that  of 
phosphoric  acid ;  with  that  of  arsenious  acid. 

3.  Experiment.  —  To  a  solution  of  arsenic  acid  (o.i  gram 
in   10  cc.  of  water)  add  magnesium  chloride  and  then  am- 
monia   until    strongly    alkaline,       Compare    with    5    under 
Phosphoric  Acid. 

4.  Experiment.  —  Add  a  little  potassium  iodide  solution 
to  some  arsenic  acid  solution,  and  warm  gently.     Is  iodine 
set  free  ?     Give  the  reaction. 

Prepare  a  faintly  alkaline  solution  of  arsenious  acid  as 
follows  :  Dissolve  a  minute  quantity  of  arsenious  oxide  in 
not  more  than  2  or  3  drops  of  hydrochloric  acid ;  dilute 
to  10  cc.  and  add,  without  heating,  a  considerable  amount 
of  sodium  bicarbonate  in  excess  of  what  is  necessary  to 
neutralize  the  acid.  To  this  solution  add,  drop  by  drop, 
a  solution  of  iodine,  and  determine  if  any  free  iodine  dis- 
appears. Write  the  reaction.  It  is,  so  far  as  the  state  of 
oxidation  of  the  arsenic  is  concerned,  exactly  the  reverse 
of  the  one  preceding.  Recall  a  previous  instance  in  which 
the  direction  of  a  reaction  of  oxidation  and  reduction  is 
changed  on  passing  from  an  acid  to  an  alkaline  solution. 

5.  Experiment.  —  Test  for  arsenious  acid  in  your  prep- 
aration:   Dissolve  a  small  crystal  of  the  arsenic  acid  in  a 
few  centimeters  of  water ;  add  sodium  bicarbonate  until  no 
more  effervescence  occurs,  and  then  a  considerable  quantity 


IO6  ELEMENTS    OF    THE    FIFTH    GROUP 

in  addition.  Add  to  this  a  solution  of  iodine,  drop  by  drop. 
The  amount  of  the  latter  which  is  decolorized  (if  any)  cor- 
responds to  the  amount  of  arsenious  acid  which  was  in  the 
sample. 


28      ANTIMONY    TRICHLORIDE    FROM    STIBNITE 
(BY-PRODUCT:  ANTIMONY  OXYCHLORIDE) 

Native  antimony  sulphide  (stibnite)  dissolves  quite  readily 
in  hydrochloric  acid,  yielding  antimony  trichloride, 

Sb2S3  +  6HC1  =  2SbCl3  +  3H2S. 

If  the  solution  so  obtained  is  distilled,  there  pass  off  at 
first  only  steam  and  hydrochloric  acid,  later  a  mixture  of 
hydrochloric  acid  and  antimony  trichloride,  and  finally  pure 
antimony  trichloride. 

The  salt  antimony  trichloride  hydrolyzes  with  a  mod- 
erate amount  of  water,  giving  a  precipitate  according  to 
the  reactions, 

SbCl3  -f  2H2O  =  SbCl(OH)2  +  2HC1, 
SbCl(OH)2  =  SbOCl  +  H2O ; 

with  more  water  a  further  hydrolysis  takes  place : 
4SbOCl  +  H20  =  Sb405Cl2  +  2HC1. 

The  product  obtained  in  this  preparation  by  mixing  the 
next  to  the  last  distillates  with  a  considerable  amount  of 
water  has  the  latter  composition.  This  compound,  how- 
ever, if  repeatedly  boiled  with  fresh  portions  of  water  may 
be  made  to  undergo  complete  hydrolysis,  leaving  finally 
only  Sb2O3. 

Procedure.  —  Treat  150  grams  of  powdered  stibnite  in 
an  8-inch  dish  at  the  hood  with  750  cc.  of  commercial 
hydrochloric  acid  (sp.  gr.  1.2)  ;  warm  the  mixture  slightly 
and  keep  it  at  50-70°,  with  frequent  stirring,  for  20  minutes 


ANTIMONY    TRICHLORIDE  IO/ 

Finally,  boil  the  solution  for  5  minutes.  Then  add  15  cc. 
more  of  concentrated  hydrochloric  acid ;  filter  the  solution 
through  asbestos  felt  (Note  4  (//))  which  has  previously  been 
moistened  with  hydrochloric  acid,  and  rinse  the  residue 
onto  the  filter  with  an  additional  15  cc.  of  hydrochloric  acid. 
Evaporate  the  filtrate  in  an  open  dish  to  200  cc. ;  then 
transfer  it  to  a  retort,  in  the  bottom  of  which  is  placed,  to 
prevent  bumping,  about  a  teaspoonful  of  small  bits  cracked 
from  an  unglazed  porcelain  dish.  Place  the  retort  on  a 
sand  bath  and  distill,  after  first  covering  the  bulb  of  the 
retort  with  an  asbestos  mantle  to  prevent  loss  of  heat.  At 
first  insert  the  neck  of  the  retort  into  a  liter  flask  half  filled 
with  cold  water  (to  absorb  the  hydrochloric  acid).  When  a 
little  of  the  distillate  begins  to  give  a  precipitate  on  dropping 
into  a  tube  of  cold  water,  exchange  the  receiving  flask  for 
a  smaller  dry  one  and  continue  the  distillation  until  a  drop 
of  the  distillate  will  solidify  when  cooled  on  a  watch  glass. 
Save  the  portion  thus  obtained  for  later  use  and  continue 
distilling,  using  a  wide  6-inch  test  tube,  which  has  previously 
been  weighed,  as  a  receiving  vessel,  until  the  liquid  is  all 
driven  out  of  the  retort.  Stopper  the  test  tube  tightly 
and  preserve  the  preparation  in  it.  If  the  product  thus 
obtained  is  not  white  it  should  be  dissolved  in  concentrated 
hydrochloric  acid  and  redistilled. 

NOTE.  —  In  case  the  stibnite  used  contains  a  considerable  quantity 
of  silicates  soluble  in  acids,  there  will  be  left  in  the  retort  as  the  distilla- 
tion progresses  a  quantity  of  gelatinous  silicic  acid  which  is  liable  to 
interfere  with  obtaining  distinct  fractions  of  the  distillate.  In  such  a 
case  distill  until  the  residue  in  the  retort  is  left  dry,  but  without  making 
the  final  change  in  receiving  vessels.  Then  pour  all  of  the  distillate 
containing  any  of  the  antimony  salt  into  a  fresh  retort  and  distill  again, 
this  time  separating  the  fractions. 

Antimony  Oxychloride.  —  Pour  the  portion  of  the  dis- 
tillate saved  from  the  above  procedure  into  2  liters  of  water. 
Stir,  allow  to  settle,  and  draw  off  the  clear  liquid.  Stir  up 
with  water  once  more,  let  settle,  and  draw  off  as  much  of 


IO8        ELEMENTS  OF  THE  FIFTH  GROUP 

the  water  as  possible.  Drain  the  precipitate  on  a  Witt 
filter,  and  spread  it  on  an  unglazed  plate  to  dry. 

Questions 

1.  Experiment.  —  Prepare   a  solution   of  antimony  tri- 
chloride.      How    can    the   formation    of    a   precipitate   be 
avoided  ?     Pass  hydrogen  sulphide  into  this  solution.     What 
is   the   precipitate  ?      How   could   it   be   converted   into   a 
product  like  stibnite  ? 

2.  How  would  the  reactions  of  phosphorus  and  arsenic 
trichlorides  with    water    differ   from    that   of   antimony  tri- 
chloride ?     Is   hydrolysis   more  or   less    complete  in   these 
cases  ?     Why,  then,  is  there  no  precipitate  ? 

29.     SODIUM    SULPHANTIMONATE    Na3SbS4.9H2O 

The  oxides  of  arsenic  and  antimony,  and  more  particu- 
larly the  higher  oxides,  are  acidic  in  nature ;  thus  with  basic 
oxides  they  will  form  salts : 

3Na2O  +  As2O8  =  2Na3AsO8; 
3Na2O  -j-  As2O5  =  2Na3AsO4. 

Sulphur,  in  accord  with  its  similarity  to  oxygen,  can  be  sub- 
stituted for  the  latter  in  many  of  its  compounds  without 
essentially  altering  their  chemical  nature,  and  the  com- 
pounds thus  obtained  have  the  same  nomenclature  as  the 
corresponding  oxygen  compounds,  except  that  the  syllable 
"  sulph  "  or  "  sulpho  "  is  inserted.  Thus  sulpho-salts  are 
produced  in  the  same  manner  as  the  oxy-salts  above: 

3Na2S  +  As2S3  —  2Na3AsS3; 
3Na2S  -|-  As2S5  —  2Na3AsS4. 

The  sulpho-salts  of  arsenic,  antimony,  and  stannic  tin  are 
particularly  characteristic  of  these  metals.  They  are  easily 
produced,  and  all  are  soluble.  They  are  stable  in  neutral 


SODIUM    SULPHANTIMONATE  IOQ 

or  basic  solutions,  but  are  decomposed  by  acids,  because  the 
anions  of  the  salts  combine  with  hydrogen  ions  to  produce 
the  very  weak  sulpho-acids,  which,  being  unstable,  decom- 
pose at  once  into  the  sulphides  of  the  metals  and  hydrogen 
sulphide : 

6H+  +  2AsS8— '  -*  2H8AsS8  -*  3H2S  +  As2S3 ; 
6H+  -f  2AsS4 -*  2H3AsS4  ->  3H2S  -j-  As2S5. 

Sodium  sulphantimonate  can  be  prepared  from  stibnite 
by  the  combined  action  of  a  solution  of  sodium  sulphide 
and  sulphur: 

28  -f-  Sb2S3  -»  Sb2S5; 
3Na2S  +  Sb2S5  -*  2Na3SbS4, 

and  it  crystallizes  well  with  nine  molecules  of  water. 

Procedure.  —  To  50  grams  of  powdered  stibnite,  107 
grams  of  crystallized  sodium  sulphide  (NasS.QHaO),1  and 
10  grams  of  powdered  sulphur  in  a  porcelain  dish,  add 
100  cc.  of  water,  bring  to  a  boil,  and  keep  at  the  boiling 
temperature  for  \  of  an  hour.  Filter  and  rinse  the  residue 
in  the  dish  and  on  the  filter  with  hot  water,  bringing  up  the 
volume  of  the  solution  to  200  cc.  While  still  hot  put  it 
away  in  a  covered  dish,  with  a  towel  placed  over  it,  to 
crystallize.  Drain  the  crystals  ;  evaporate  the  mother  liquor 
somewhat  to  obtain  a  second  crop  of  crystals.  Spread  the 
crystals  on  a  porous  plate,  and  stopper  them  tightly  as  soon 

as  dry. 

Questions 

i.  Experiment. —  Prepare  a  little  precipitated  antimo- 
nous  sulphide.  How?  Treat  this  precipitate  with  a  solu- 
tion of  ammonium  polysulphide,  (NH4)2SX.  Discuss,  with 
reactions,  the  nature  of  the  soluble  compound  produced. 
Finally,  acidify  the  solution  with  hydrochloric  acid.  What 

is  the  reaction? 

• 

1  Or  use  35  grams  of  anhydrous  sodium  sulphide  and  an  additional 
72  cc.  of  water. 


IIO  ELEMENTS    OF    THE    FIFTH    GROUP 

30.     ANTIMONY    PENTASULPHIDE,  Sb2S5 

This  compound  cannot  be  prepared  directly  from  the 
trisulphide  and  sulphur,  because  it  is  decomposed  at  a 
temperature  below  that  at  which  the  latter  substances 
would  react.  As  has  just*  been  seen,  however,  the  higher 
sulpho-salt  of  antimony  can  be  readily  prepared  in  the  wet 
way ;  and  this,  on  decomposition  with  a  dilute  acid,  yields 
antimony  pentasulphide.  This  substance  is  much  used  in 
vulcanizing  rubber. 

Procedure.  —  Dissolve  40  grams  of  the  sodium  sulphan- 
timonate  obtained  in  the  last  preparation,  and  dilute  with 
i  liter  of  cold  water.  Mix  15  cc.  of  concentrated  sulphuric 
acid  with  350  cc.  of  water,  cool,  and  place  in  a  2-liter,  or, 
better,  a  3  or  4-liter  common  bottle.  To  this  add  slowly, 
and  with  constant  stirring,  the  solution  prepared  above. 
Fill  the  bottle  with  water  and  stir  thoroughly.  Let  the 
precipitate  settle,  draw  off  the  liquid,  and  wash  by  decan- 
tation  until  the  wash  water  no  longer  gives,  with  barium 
chloride,  the  test  for  a  sulphate.  After  the  last  washing 
let  settle  for  some  time,  draw  off  as  much  as  possible  of 
the  clear  liquid,  and  transfer  the  slime  to  a  large  filter 
(Note  4  (c] ;  do  not  omit  to  reenforce  the  point  of  the 
filter)  to  drain  for  12  hours  or  longer.  Without  removing 
the  pasty  antimony  sulphide  from  the  filter,  open  out  the 
latter  on  an  unglazed  plate,  and  leave  it  on  a  shelf  above 
the  steam  table  where  the  temperature  does  not  rise  above 
50°.  When  the  product  is  completely  dry,  detach  the 
hardened  lumps  from  the  paper  and  pulverize  them  in  a 
mortar. 

Questions 

i.  Write  all  the  reactions  involved  in  the  preparation 
of  antimony  pentasulphide  from  stibnite. 


METALLIC    ANTIMONY  I  I  I 

31.     METALLIC    ANTIMONY 

This  metal  is  obtained  on  a  commercial  scale  both  by 
reducing  antimony  oxide  with  carbon  and  by  reducing 
antimony  sulphide  by  means  of  metallic  iron.  The  latter 
method  possesses  the  advantage  that  antimony  sulphide, 
a  natural  product,  is  used  directly  and  does  not  need  to 
be  first  converted  into  the  oxide.  The  iron  sulphide  formed 
by  this  method  is  fusible  and  forms  a  slag ;  but  the  slag  is 
made  more  fusible  by  the  admixture  of  sodium  sulphate  as 
directed,  and  thus  the  globules  of  melted  antimony  are 
allowed  to  sink  more  easily  to  the  bottom  of  the  crucible 
and  form  a  metallic  regulus.  The  slag  furthermore  covers 
the  surface  of  the  metal  and  hinders  its  volatilization  and 
oxidation. 

Procedure.  —  Mix  100  grams  of  stibnite  with  42  grams 
of  iron  filings,  10  grams  of  anhydrous  sodium  sulphate,  and 
2  grams  of  powdered  charcoal,  and  place  the  mixture  in  a 
clay  crucible.  Cover  the  crucible  tightly,  and  heat  it  in 
the  gas  furnace  for  one  hour  at  a  bright  red  heat.  The 
temperature  should  not  be  high  enough  to  volatilize  the 
antimony,  which  would  in  that  case  escape  as  a  white 
smoke  consisting  of  antimony  oxide,  yet  the  slag  of  iron 
sulphide  must  be  completely  softened,  although  it  should 
not  melt  to  a  thin  liquid.  After  about  half  an  hour  test 
the  conditions  by  removing  the  cover  a  moment  and  stirring 
the  slag  with  an  iron  rod  to  see  whether  it  is  in  the  proper 
semi-fluid  condition.  After  the  reaction  is  complete,  allow 
the  crucible  to  cool,  break  it  and  separate  the  regulus  of 
antimony  from  the  slag. 

Questions 

i.  Experiment.  —  Warm  a  piece  of  metallic  antimony 
with  hydrochloric  acid.  Where  does  antimony  stand  in 
the  electromotive  series? 


112  ELEMENTS    OF    THE    FIFTH    GROUP 

2.  Experiment.  —  Warm  a  little  powdered  antimony 
with  nitric  acid  in  a  casserole.  What  is  the  product? 
Compare  it  with  the  product  obtained  by  treating  tin  in 
the  same  manner. 


32.      BISMUTH    BASIC    NITRATE   (BISMUTH 
SUBNITRATE) 

Although  bismuth  is  the  most  strongly  metallic  element 
of  the  fifth  group,  yet  its  salts  in  aqueous  solution  undergo 
partial  hydrolysis  very  readily.  In  presence  of  a  consider- 
able amount  of  free  acid,  the  Bi+  +  +-ion  is  capable  of  exist- 
ence in  solution;  but  with  decreasing  quantities  of  acid 
the  tendency  of  water  to  produce  hydrolysis  increases,  and 
the  basic  salt  of  bismuth,  which  is  only  slightly  soluble, 
separates  : 


OH  _}-  2H++  2NO,f. 


On  pouring  a  solution  of  bismuth  nitrate  into  a  considerable 
quantity  of  cold  water  the  basic  nitrate,  according  to  the 
above  formula,  is  precipitated.  This  salt,  however,  is  not 
stable  in  contact  with  a  solution  which  does  not  contain 
nitric  acid  of  a  concentration  of  at  least  about  \  molal,  but 
slowly  changes  over  into  some  other  more  basic  nitrate, 
and  if  washed  repeatedly  with  pure  water  will  finally  go 
over  completely  into  the  hydroxide  : 


Bi  —  OH  +  H2O  ^  Bi  —  OH  -f  H+  +  NO-T  . 


Under  the  conditions  given  in  the  following  procedure,  this 
production  of  a  more  basic  salt  will  occur  if  the  precipitate 
is  allowed  to  stand  in  contact  with  the~solution  for  a  con- 
siderable time;  hence  the  directions  to  filter  at  once. 


BISMUTH    BASIC    NITRATE  I  I  3 

The  basic  nitrate  is  by  no  means  completely  insoluble 
in  water,  and  the  filtrate  contains  considerable  quantities 
of  bismuth,  which  can  be  conveniently  saved  as  oxide  by 
precipitating  with  sodium  carbonate. 

Procedure.  —  Dissolve  without  heating  40  grams  of  crys- 
tallized bismuth  nitrate,  Bi(NO8)3.5H2O,  in  10  cc.  of  nitric 
acid  (sp.  gr.  1.2)  and  20  cc.  of  water.  Pour  this  into 
2  liters  of  cold  water  and  stir  thoroughly  for  a  few  minutes. 
Let  the  precipitate  settle  completely,  and  as  soon  as  this 
has  occurred  draw  off  and  save  the  supernatant  liquor; 
drain  the  precipitate  on  a  suction  filter,  and  wash  it  quickly 
with  about  20  cc.  of  water.  Dry  the  precipitate  at  the 
steam  table,  and  preserve  it  as  a  powder. 

Bismuth  Oxide.  —  Combine  all  the  liquors  from  the  fore- 
going; add  sodium  carbonate  until  alkaline  to  litmus;  let 
settle,  and  draw  off  the  supernatant  liquor ;  boil  the  remain- 
ing suspension  after  adding  to  it  about  20  grams  more  of 
sodium  carbonate.  Then  wash  the  precipitate  twice  by 
decantation,  drain  on  a  suction  filter,  and  wash  with  two 
or  three  portions  of  water.  Dry  and  preserve. 

Questions 

1.  In    accordance   with   the   above   directions,  sodium 
carbonate  is  used  to  precipitate  bismuth  hydroxide.     Why 
should  not  the  precipitate  be  bismuth  carbonate  ? 

2.  If  this  precipitate  is  not  finally  boiled  with  an  excess 
of  sodium  carbonate,  it  is  likely  to  contain  a  certain  amount 
of  basic  nitrate.     Explain  why  this  should  be  so  and  why 
the  boiling  will  convert  it  completely  into  the  hydroxide. 


114  ELEMENTS    OF    THE    FIFTH    GROUP 

GENERAL    QUESTIONS.     V 

ELEMENTS   OF   THE   FIFTH    GROUP    OF   THE   PERIODIC    SYSTEM 

Experiments 

(The  results  observed  are  to  be  recorded  in  the  laboratory  note- 
book at  the  time  the  experiments  are  performed.) 

1.  Boil  about  \  gram  of  powdered  metallic  arsenic  with 
nitric  acid  (1.2)  until  the  metal  is  entirely  dissolved.     Evap- 
orate the  solution  just  to  dryness  by  heating  it  over  a  free 
flame  in  a  casserole  while  holding  the  latter  in  the  hand  and 
rotating  its  contents.     In  this  way  all  the  unused  nitric  acid 
is  expelled.     When  cool  add  10  cc.  of  water  and  warm  until 
the  arsenic  acid  has  all  dissolved.     Prove  that  the  solution 
contains  an  acid  (/.  <?.,  more  than  an  accidental  trace  of  nitric 
acid,  which  might  not   have  been  completely  expelled  and 
might  be  sufficient  to  redden  litmus). 

Boil  about  \  gram  of  powdered  antimony  with  nitric 
acid  (1.2).  What  is  the  character  of  the  product  formed? 

Treat  \  gram  of  powdered  bismuth  in  the  same  manner. 
Evaporate  the  solution  so  obtained  to  a  very  small  volume 
(about  i  cc.)  to  remove  most  of  the  surplus  nitric  acid. 
Then  pour  the  solution  into  50  cc.  or  more  of  cold  water. 

2.  Pass  hydrogen  sulphide  into  hot  dilute  solutions  of 
arsenic,  antimony,  and  bismuth  trichlorides  in  separate  test 
tubes.     Note  the  color  of  the  precipitates.      Collect  each 
precipitate  on  a  filter  and  treat  it  with  warm  ammonium 
polysulphide  (ammonium  sulphide    in  which  more   sulphur 
is  dissolved).     Lastly,  acidify  with  hydrochloric  acid  each 
of  the  solutions  so  obtained. 

Questions 

I.  Of  Group  V  only  those  elements  falling  in  the  right-hand 
sub-column  need  especial  consideration,  namely,  nitrogen,  phosphorus, 
arsenic,  antimony,  and  bismuth.  In  this  series  we  pass  by  a  very 


GENERAL    QUESTIONS    V  I  I  5 

perfect  gradation  from  nitrogen,  a  pronounced  non-metal,  to  bismuth, 
a  quite  pronounced  metal. 

State  in  each  case  whether  the  nitrate  or  sulphate  of 
trivalent  phosphorus,  arsenic,  antimony,  or  bismuth  can  be 
prepared,  and  if  so  whether  it  can  be  dissolved  in  water 
without  suffering  complete  hydrolysis  (see  Experiment  i). 
How  does  the  basic  nature  of  the  trioxide  change  in  the 
series,  phosphorus  to  bismuth  ?  Can  nitrates  or  sulphates 
of  any  of  these  elements  in  their  pentavalent  condition  be 
prepared  ?  For  any  one  of  the  elements,  which  is  the  more 
strongly  basic  in  nature,  the  trioxide  or  the  pentoxide  ? 
Which  is  the  more  strongly  acidic  ?  Give  the  symbol  of 
the  most  common  acid,  if  one  exists,  which  is  derived  from 
the  pentoxide  of  each  of  these  elements.  How  does  the 
acidic  nature  of  the  pentoxide  change  in  passing  from 
nitrogen  to  bismuth  ? 

2.  Name  the  simplest  hydrogen  compound  of  each  of 
the    first   four    elements.     Compare   the    stability   of    these 
hydrides    when    heated.      Compare    any    ability   they   may 
possess  to  unite  with  water  to  form  bases,  and  with  acids 
to  form  salts. 

3.  Write  the  equations  for  the  reaction  of  nitric  acid 
with  phosphorus,   arsenic,   antimony,   and  bismuth,   respec- 
tively (see  Experiment  i). 

4.  How  do  the  trisulphides   of   arsenic  and   antimony 
behave  when  treated  with  a  solution  of  ammonium  sulphide, 
(NH4)2S  ?      With    a    solution    of    ammonium    poly  sulphide, 
(NH4)2S.SX?     How    does   the    solution   obtained   with    am- 
monium  polysulphide  behave  when   it  is  acidified  ?     Give 
equations  for  all  the  reactions.     (See  Experiment  2.) 

What  is  the  relation  between  sulpho-  and  oxy-acids  ? 
Show,  for  example,  how  sodium  sulpharsenate  is  derived 
from  two  simple  sulphides,  and  sodium  arsenate  from  the 
corresponding  oxides. 


CHAPTER   VI 

HEAVY    METALS    OF    THE    SIXTH,    SEVENTH, 

AND    EIGHTH    GROUPS    OF   THE 

PERIODIC    SYSTEM 

By  turning  to  the  Periodic  Table  of  the  elements  it  is 
observed  that  chromium,  manganese,  iron,  cobalt,  and  nickel, 
and  following  these  copper  and  zinc,  come  in  the  middle 
portion  of  the  long  period  that  begins  with  potassium  and 
ends  with  bromine.  The  seven  elements  mentioned  possess 
high  specific  gravities  and  all  come  under  the  classification 
of  heavy  metals.  In  certain  of  their  compounds  they  are 
extremely  similar  to  one  another ;  in  other  of  their  properties 
they  are  very  dissimilar  and  exhibit  the  chemical  character- 
istics of  the  respective  groups  to  which  they  belong. 

The  heavy  metals  occupying  a  corresponding  position 
in  the  middle  of  the  next  long  period  are  molybdenum,  an 
unknown  element  which  should  come  below  manganese, 
ruthenium,  rhodium,  palladium,  silver,  and  cadmium.  In 
the  middle  of  the  next  long  period  come  tungsten,  another 
unknown  element  which  should  resemble  manganese,  os- 
mium, iridium,  platinum,  gold,  and  mercury.  In  the  last 
long  period,  of  which  there  is  at  best  only  a  fragmentary 
indication,  the  only  representative  of  this  class  of  heavy 
metals  is  uranium. 

In  the  sixth  group,  chromium,  molybdenum,  tungsten, 
and  uranium  constitute  Family  A.  In  their  trioxides  they 
show  the  characteristic  valence  of  the  sixth  group  and 
resemble  in  properties  the  non-metals  of  Family  B,  of  which 
sulphur  is  the  type.  In  their  lower  oxides  they  show  none 

117 


Il8  HEAVY     METALS 

of  the  group  characteristics,  but  show  the  general  base- 
forming  properties  of  the  heavy  metals. 

In  the  seventh  group,  manganese,  the  only  known  rep- 
resentative of  Family  A,  resembles  the  halogens  in  its 
heptoxide,  Mn2O7 ;  in  its  lower  oxides  it  shows  no  resem- 
blance to  the  halogens,  but  does  show  properties  similar  to 
those  of  other  heavy  metals  when  they  are  in  the  same  state 
of  oxidation ;  in  its  lowest  oxide,  MnO,  it  is  distinctly  a 
base-forming  element. 

In  the  eighth  group,  each  position  instead  of  being  filled 
by  a  single  element  is  occupied  by  a  group  of  three  ele- 
ments. Thus  there  appear  in  triads :  iron,  cobalt,  and 
nickel ;  ruthenium,  rhodium,  and  palladium ;  and  osmium, 
iridium,  and  platinum.  In  this  group  there  is  no  subdivi- 
sion into  A  and  B  Families,  but  all  of  the  members  are 
heavy  metals.  The  Zero  Group,  which  comprises  the  inert 
gases,  helium,  neon,  argon,  krypton,  and  xenon,  may  be 
regarded  as  bearing  the  same  relation  to  Group  VIII  as  is 
shown  in  the  other  groups  between  the  A  and  B  Families. 
If  this  view  is  a  correct  one,  the  divergence  in  properties 
between  the  families  is  in  this  case  at  a  maximum. 

Of  the  heavy  metals  discussed  above,  the  only  ones  that 
are  of  frequent  occurrence  and  that  are  to  receive  detailed 
treatment  in  this  chapter  are  chromium,  manganese,  and 
iron. 


33.     POTASSIUM   BICHROMATE   FROM    CHROMITE 

The  most  important  source  of  chromium  is  the  mineral 
chromite,  FeO.Cr2O3  or  Fe(CrO2)2.  This  substance,  as  in- 
dicated by  the  formula,  may  be  regarded  as  a  mixture  of 
ferrous  oxide  and  chromic  oxide,  or  as  a  salt,  chromite 
of  iron,  in  which  ferrous  oxide  is  the  basic  constituent 
and  chromic  oxide  the  acidic.  Chromite  is  a  difficult 
material  to  decompose,  and  the  ordinary  method  by 
which  this  is  accomplished  is  by  treatment  at  a  high 
temperature  with  an  alkali  and  an  oxidizing  agent.  By 
this  treatment  the  iron  of  the  chromite  is  converted  to 
the  ferric  condition  (Fe2O3),  and  the  chromium  is  oxi- 
dized to  the  hexavalent  condition  (CrO3),  at  the  same  time 
combining  with  the  alkali  to  form  a  chromate  (for  example, 
K2CrO4). 

In  the  commercial  method  for  manufacturing  chromates, 
atmospheric  oxygen  is  utilized  as  the  oxidizing  agent.  The 
chromite  is  mixed  with  potassium  carbonate  and  calcium 
carbonate,  the  latter  to  give  porosity,  and  then  heated  for 
a  considerable  time  in  a  furnace  with  free  access  of  air. 
The  chromium  trioxide,  CrO3,  produced  by  the  oxidation 
reacts  with  the  potassium  carbonate,  displacing  carbon 
dioxide  and  giving  potassium  chromate,  K2CrO4.  After 
cooling,  the  contents  of  the  furnace  are  treated  'with  a 
solution  of  sodium  sulphate ;  the  potassium  chromate  dis- 
solves, the  iron  oxide  is  insoluble,  and  the  calcium  oxide 
(from  heating  the  carbonate)  reacts  with  the  sodium  sul- 
phate to  form  insoluble  calcium  sulphate.  From  the 
solution  potassium  chromate  could  be  crystallized  but 
for  the  fact  that  it  is  very  soluble  in  water  and  could 
not  be  separated  thus  from  other  salts  in  the  solution. 
Potassium  bichromate,  however,  is  much  less  soluble,  and 

119 


120 


HEAVY     METALS 


if   this  is  formed    by    adding    a    sufficient    amount  of   sul- 
phuric  acid, 

2K2CrO4  +  H2SO4  =  K2SO4  +  K2Cr2O7  -f  H2O, 

it  can  be  obtained  pure  by  crystallization. 

On  account  of  .the  difficulty  of  carrying  out  the  above 
process  on  a  laboratory  scale,  the  following  less  economical 
procedure  is  given  in  which  the  mineral  is  heated  with  con- 
siderably more  than  the  quantity  of  potassium  carbonate 
theoretically  required,  in  order  to  give  a  more  liquid  melt, 
and  with  potassium  nitrate  for  the  oxidizing .  agent.  The 
solution  obtained  by  extracting  this  melt  with  water  con- 
tains so  much  potassium  carbonate  that  it  would  be  very 
difficult  to  separate  the  also  very  soluble  potassium  chro- 
mate  from  it  by  crystallization.  If,  however,  acetic  acid  is 
added  until  the  solution  reacts  slightly  acid,  the  potassium 
carbonate  is  converted  into  the  very  soluble  acetate,  and  the 
chromate  is  changed  to  the  only  moderately  soluble  potas- 
sium bichromate  which,  especially  in  the  presence  of  the 
large  amount  of  the  other  salt  with  the  K+-ion  in  common, 
can  be  very  readily  crystallized  out. 


A  saturated  solution  contains  for  each  100  grams  of  water  the  given 
number  of  grams  of  the  anhydrous  salt. 


. 

0° 

10° 

20° 

30° 

40° 

50° 

70° 

100° 

K2CrO4   

59 

61 

63 

65 

67 

69 

73 

79 

K2Cr2O7      

5 

7 

12 

20 

26 

35 

55 

88 

Procedure.  —  Mix  40  grams  of  finely  powdered  chromite1 
with  100  grams  of  potassium  carbonate  and  30  grams  of 

1  Unless  an  ore  can  be  obtained  which  approximates  the  composi- 
tion of  the  pure  mineral,  FeO.Cr2O3,  it  would  prove  more  satisfactory 
to  take  27  grams  of  pure  chromic  oxide,  Cr2C>3,  instead  of  40  grams 
of  chromite. 


POTASSIUM    BICHROMATE  121 

potassium  nitrate.  Place  the  mixture  in  a  cast-iron  crucible, 
which  it  must  on  no  account  fill  more  than  two-thirds  full, 
else  when  melted  it  will  run  over.  Heat  in  a  gas  furnace  to 
a  white  heat  (but  using  care  not  to  reach  the  very  highest 
heat,  which  might  melt  the  crucible)  until  the  melted  charge 
has  ceased  to  effervesce.  Pour  the  molten  mass  out  onto 
a  dry1  iron  plate.  When  cool  crack  it  up  and  dissolve  it, 
together  with  what  still  adheres  to  the  crucible,  in  boiling 
water.  Filter  the  solution,  and  extract  the  residue  with  a 
little  more  boiling  water  and  pour  through  the  same  filter. 
Add  glacial  acetic  acid  (cautiously)  to  the  filtrate  until  it 
has  become  acid.  Boil  down  the  solution  to  300  cc.,  or  to 
even  a  less  volume  if  no  solid  salt  begins  to  separate.  Add 
25  cc.  more  of  glacial  acetic  acid,  let  stand  for  some  time, 
and  finally  cool  to  o°  before  separating  the  crystal  meal  of 
potassium  bichromate  from  the  mother  liquor.  Purify  the 
product  by  recrystallization. 

Questions 

1.  Mention    at     least    three    oxidizing    agents    which 
might  have  been  used  instead  of  potassium  nitrate  in  this 
preparation. 

2.  How    might   the    oxidation    of    chromic    hydroxide, 
Cr(OH)8,  be  accomplished  in  the  wet  way?     Experiment. — 
To  5   cc.  of  a  cold   solution  of   a  chromic   salt  add  about 
i    gram   of   sodium    peroxide,   agitate   for  a  few  moments, 
and  then  warm  until  effervescence  ceases.     Formulate  the 
equations  for  the   intermediate  reactions  in  such  a  way  as 
to  show  the  state  of  oxidation  of    chromium  in  each    com- 
pound involved,  and  then  add  the  separate  equations  to  give 
one  for  the  complete  reaction. 

3.  Experiment.  —  To  a    solution   of   potassium  bichro- 
1  See  footnote  2,  page  60. 


122  HEAVY    METALS 

mate  add  potassium  carbonate  until  no  more  effervescence 
occurs.  Observe  and  explain  any  change  in  color.  To  a 
solution  of  potassium  chromate  add  an  acid  and  observe  as 
before.  Explain  fully  the  relation  between  chromates  and 
bichromates. 


34.     POTASSIUM   CHROMATE   FROM   POTASSIUM 
BICHROMATE 

Dissolve  50  grams  of  potassium  bichromate  in  water  and 
add  the  calculated  amount  of  potassium  carbonate  dissolved 
in  water.  The  color  should  just  change  to  clear  yellow,  and 
no  trace  should  be  left  of  the  reddish  hue  characteristic  of 
the  bichromate.  Crystallize  the  product  from  the  solution 
(see  solubility  table  on  page  120). 

Answer  the  questions  given  under  Potassium  Bichromate. 

35.     CHROMIC   ANHYDRIDE,    CrO8 

When  a  chromate  or  a  bichromate  is  treated  with  a 
strong  acid,  chromic  acid  is  formed  in  the  solution.  The 
affinity  of  chromic  anhydride  for  water  is  far  less  than  that 
of  sulphuric  anhydride  for  water ;  and  chromic  acid,  there- 
fore, instead  of  existing  in  solution  entirely  in  the  form 
H2CrO4,  is  broken  down  to  a  great  extent  into  H2Cr2O7 
(i.e.,  H2O.2CrO3)  and  even  to  CrO8.  Especially  in  the 
presence  of  a  large  amount  of  sulphuric  acid,  the  last  form 
is  produced  so  freely  that  it  crystallizes  out  in  the  shape 
of  red  needles. 

Commercially,  chromic  anhydride  is  most  often  prepared 
by  the  action  of  sulphuric  acid  on  potassium  bichromate. 
Potassium  acid  sulphate  is  first  crystallized  from  the  mixture 
and  after  that  the  chromic  anhydride;  but  a  good  deal  of 
care  is  necessary  to  obtain  the  product  uncontaminated  with 


CHROMIC    ANHYDRIDE  123 

potassium  salt.  When,  however,  lead  chromate  is  used  as 
the  source  of  the  chromic  acid,  the  lead  can  be  completely 
removed,  since  with  sulphuric  acid  the  extremely  insoluble 
lead  sulphate  is  formed.  The  remaining  solution  then  con- 
tains nothing  but  chromic  acid  and  an  excess  of  sulphuric 
acid.  The  chromic  anhydride  can  then  be  crystallized  out, 
it  being  least  soluble  in  a  mixture  containing  in  the  neigh- 
borhood of  75  per  cent,  of  sulphuric  acid. 

Procedure.  —  Lead  Chromate.  Dissolve  100  grams  of 
lead  acetate  in  i  liter  of  water,  and  add  a  few  drops  of 
acetic  acid  if  necessary  to  clear  up  any  turbidity.  Dissolve 
39  grams  of  potassium  bichromate  in  i  liter  of  water,  and 
add  this  solution  to  the  first,  while  stirring.  Wash  the  pre- 
cipitate by  decantation  until  less  than  0.5  per  cent,  of  the 
soluble  salt  remains  (Note  5  (b)  on  page  13);  then  collect 
the  lead  chromate  on  an  ordinary  filter  (Note  4  (<:)),  and  after 
draining  dry  it  thoroughly. 

Chromic  Anhydride.  —  Take  the  lead  chromate  prepared 
above  (or  100  grams  of  a  commercial  sample),  reduce  all 
lumps  to  a  fine  powder,  add  200  grams  of  concentrated  sul- 
phuric acid  (in  cc.),  and  stir  the  mixture  with  a  pestle  until 
a  perfectly  smooth  paste  is  produced.  Allow  the  mixture  to 
digest  24  hours  in  a  warm  place,  as  on  the  shelf  above  the 
steam  table.  Dilute  to  a  liter  with  water,  and  filter  the  solu- 
tion through  asbestos  felt  (Note  4  (d)  on  page  10).  Wash 
the  lead  sulphate  until  it  is  nearly  or  quite  white,  separate 
it  from  the  asbestos  as  completely  as  possible,  and  preserve 
it  as  a  by-product.  Evaporate  the  filtrate  in  a  porcelain 
dish  until  crystals  of  chromic  anhydride  begin  to  form  a 
scum  on  the  surface  of  the  liquid.  Let  the  solution  cool 
slowly ;  then  collect  the  crystals  in  a  funnel  in  which  a  per- 
forated plate  or  a  glass  marble  is  placed,  and  drain  out  all 
the  liquid  with  suction.  Cover  the  funnel  with  a  watch 
glass  while  evaporating  the  liquor  to  obtain  a  second  crop 


124  HEAVY    METALS 

of  crystals.  Collect  this  crop  in  the  same  funnel  together 
with  the  first  crop,  and  wash  the  product  twice  with  con- 
centrated nitric  acid  to  remove  the  sulphuric  acid  adhering. 
Use  sufficient  nitric  acid  each  time  to  just  wet  the  entire 
mass  of  crystals,  and  then  drain  it  off  as  thoroughly  as  pos- 
sible with  the  suction  (Note  5  (a)  on  page  12).  Finally, 
drive  off  the  nitric  acid  by  heating  the  crystals  very  cau- 
tiously in  a  small  porcelain  dish  placed  on  a  sand  bath. 
Keep  turning  over  the  mass  of  crystals  with  a  glass  spatula 
to  avoid  heating  the  lower  layer  too  strongly.  Chromic 
anhydride  melts  at  192°,  and  care  must  be  taken  to  avoid 
reaching  this  temperature.  When  the  crystals  appear  dry 
and  no  more  white  vapor  can  be  detected  on  breathing 
across  them,  preserve  them  in  a  glass-stoppered  bottle. 

Questions 

1.  When  chromic  anhydride  is  dissolved  in  water,  what 
components   are  produced  in  the  solution  ?     What  salt  is 
precipitated  if  lead   acetate  or  barium  acetate  is  added  to 
this    solution?      Why    is    it    not   the    bichromate    which    is 
obtained  ? 

2.  Experiment.  —  Heat    a    little     chromic     anhydride 
strongly  on  a  bit  of  porcelain.     What  color  change  occurs  ? 
(The  color  of  the  product  can  be  better  observed  if  a  par- 
ticle  is   pulverized   in    a  white   mortar.)      Is   the   product 
soluble  in  water  ?     In    hydrochloric   acid  ?     What   relation 
does  it  bear  to  the  mineral  chromite? 

36.     CHROMIC    ALUM 

The  preparation  of  potassium  bichromate  illustrated  how 
chromic  oxide,  Cr2O3,  as  it  exists  in  nature  as  a  constituent 
of  the  mineral  chromite,  can  be  oxidized  to  a  chromate  in 
which  chromium  exists  as  CrO3.  For  the  preparation  of 


CHROMIC    ALUM  125 

chromic  alum,  K2SO4.Cr2(SO^)B.24H2O,  it  might  seem  as 
if  chromite  should  yield  chromic  sulphate  directly  on  treat- 
ment with  sulphuric  acid.  This  is,  however,  impossible, 
because  the  natural  material  is,  as  already  stated,  very 
resistant  to  the  action  of  acids,.  It  yields  only  to  the  action 
of  powerful  oxidizing  agents,  which  convert  it  into  a  chro- 
mate,  and  therefore  potassium,  or  sodium,  bichromates  are 
always  the  products  made  directly  from  the  mineral,  and 
these  serve  as  the  materials  from  which  other  compounds 
of  chromium  are  prepared.  To  make  chromic  alum  from 
potassium  bichromate  it  is  necessary  to  reduce  the  chro- 
mium to  the  same  state  of  oxidation  in  which  it  originally 
existed  in  the  mineral,  and  to  add  sufficient  sulphuric  acid 
to  form  the  sulphates  of  potassium  and  chromium.  Alcohol 
may  be  used  as  the  reducing  agent,  it  being  itself  oxidized 
to  aldehyde,  a  body  whose  presence  is  made  very  evident 
by  its  penetrating  odor. 

Chromic  alum  is  isomorphous  with  common  alum  and 
can  easily  be  obtained  in  large  and  beautiful  deep  purple 
crystals.  Care  must,  however,  be  exercised  not  to  allow 
the  temperature  of  its  solution  to  rise  above  50°  during 
the  preparation,  for  when  heated  beyond  this  point  it  under- 
goes a  change  into  a  green  noncrystallizable  body.  This 
green  body  is  not  stable  at  the  ordinary  temperature,  and 
after  cooling  it  will  change  slowly  back  into  the  ordinary 
crystallizable  chromic  alum ;  but  so  slowly,  however,  that  if 
once  it  is  formed  the  preparation  is  practically  spoiled. 

At  25°,  24  grams  of  K2SO4.Cr2(SO4)3.24H2O  will  dissolve 
in  100  grams  of  water,  and  the  solubility  increases  very 
rapidly  with  the  temperature. 

Procedure.  —  Pulverize  100  grams  of  potassium  bichro- 
mate, and  cover  it  in  an  8-inch  evaporating  dish  with 
400  cc.  of  water.  Add  78  cc.  of  concentrated  sulphuric 
acid,  and  stir  until  the  bichromate  is  all  dissolved.  Adding 


126  HEAVY    METALS 

the  sulphuric  acid  should  produce  enough  heat  to  dissolve 
the  bichromate,  but  if  it  is  necessary  heat  the  mixture  a 
little  more.  Be  sure  that  the  last  trace  of  solid  is  dissolved. 
Allow  the  solution  to  cool  to  40° ;  then  add  alcohol,  a  drop 
at  a  time,  while  stirring  constantly  with  the  stem  of  a  ther- 
mometer until  the  temperature  commences  to  rise.  Then 
place  the  dish  in  a  pan  of  ice  and  water  and  add  alcohol, 
65  cc.  in  all,  at  first  very  slowly,  endeavoring  to  keep  the 
temperature  between  35°  and  40°,  and  finally  more  rapidly. 
Keep  the  temperature  at  all  times  well  below  50°;  and  if  it 
should  start  to  rise  suddenly,  due  to  too  large  an  addition 
of  alcohol,  and  get  as  high  as  50°,  drop  a  piece  of  ice 
directly  into  the  solution.  Finally,  let  the  solution  cool 
completely  in  the  bath  of  ice  water,  or,  still  better,  let  it 
stand  over  night.  Collect  the  crystal  meal  on  a  Witt  filter 
and  suck  it  free  from  liquid.  Recrystallize  so  as  to  obtain 
large,  well-shaped  crystals,  following  a  similar  procedure  and 
observing  the  same  precautions  as  with  common  alum  (see 
page  54).  A  saturated  solution  of  this  salt  should  be  .pre- 
pared at  35°.  After  freeing  it  of  any  undissolved  particles 
of  the  crystal  meal,  warm  it  to  40°,  and  set  it  to  crystallize, 
with  the  addition  of  about  ten  very  small  crystals  to  serve 
as  nuclei.  Dry  with  filter  paper  the  crystals  so  obtained, 
and  stopper  them  at  once  in  a  bottle,  since  they  are  quite 
efflorescent. 

Questions 

1.  Formulate  the  equations  for  the  separate  reactions 
involved  in  the  reduction  of  the  bichromate,  in  such  a  way 
as  to   show  the  changes  occurring  in   the  state  of   oxida- 
tion of  the  chromium ;   the  alcohol,  C2H6O,  is  oxidized  to 
aldehyde,  C2H4O. 

2.  Sulphur  dioxide  might  serve  as  the  reducing  agent. 
Give  equations  for  the  partial  and  complete  reactions. 


CHROMIUM  127 

3.  Dissolve  -J-  gram  of  potassium  bichromate  in  10  cc. 
of  water  and  add  10  cc.  of  dilute  sulphuric  acid.  Heat  to 
boiling,  and  pass  in  hydrogen  sulphide  until  the  color  is 
changed  completely  to  green.  To  what  is  the  green  color 
due  ?  What  is  the  precipitate  ?  Formulate  equations  also 
for  this  reaction. 


37.     CHROMIUM    METAL   BY    THE    GOLDSCHMIDT 
PROCESS 

The  readiest  method  of  obtaining  the  metal  chromium 
from  its  oxide,  and  one  which  yields  it  in  a  high  state  of 
purity,  is  the  so-called  Goldschmidt  Process,  in  which  use  is 
made  of  metallic  aluminum  as  the  reducing  agent  according 
to  the  reaction, 

2A1  +  Cr2O3  =  A12O3  -f-  2Cr. 

The  heat  produced  by  the  oxidation  of  aluminum  is  so  great 
that  it  is  sufficient  to  effect  the  decomposition  of  the  chromic 
oxide  with  still  enough  surplus  heat  to  produce  a  tempera- 
ture high  enough  to  melt  the  metallic  chromium.  It  is 
evident  that  before  this  reaction  can  be  made  to  progress 
spontaneously  a  sufficient  temperature  must  be  developed 
to  decompose  the  chromium  oxide.  This  necessary  tem- 
perature is  a  good  deal  higher  than  that  of  a  Bunsen  flame 
or  of  a  common  furnace,  but  can  be  obtained  by  use  of  the 
fuse  powder  described  below.  When  once  started  in  this 
way  the  reaction  itself  produces  a  temperature  high  enough 
to  insure  its  continuance. 

Carried  out  on  the  small  scale  of  a  laboratory  prepa- 
ration, the  heat  produced  is  not  quite  sufficient  to  melt  the 
metal  and  slag  so  thoroughly  that  the  metal  can  settle  out 
to  form  a  compact  regulus  at  the  bottom  of  the  crucible. 
By  adding  a  small  amount  of  potassium  bichromate  to  the 


128  HEAVY    METALS 

charge,    however,    the    reaction    becomes    more    energetic, 
owing  to  the  more  available  supply  of  oxygen. 

Procedure.  —  Heat  some  powdered  chromic  oxide  in  an 
iron  pan  over  a  Bunsen  burner.  Melt  some  potassium 
bichromate  in  a  clean  iron  pan  and  pulverize  it  in  a  mortar 
after  it  has  again  solidified.  It  is  necessary  for  the  mate- 
rials used  to  be  entirely  free  from  moisture.  Mix  210  grams 
of  the  ignited  chromic  oxide,  60  grams  of  the  fused  potas- 
sium bichromate,  and  96  grains  of  granulated  aluminum 
(not  the  powder  which  is  used  for  a  pigment),  pack  the 
mixture  closely  into  a  clay  crucible,  and  embed  the  latter 
in  a  pail  of  sand.  Make  a  hole  about  4  cm.  deep  in  the 
middle  of  the  charge,  and  fill  it  with  about  10  grams  of  a 
fuse  powder  made  from  10  parts  of  barium  peroxide  and 
i  part  of  granulated  aluminum.  Insert  a  strip  of  mag- 
nesium ribbon  into  the  fuse  powder.  Place  the  whole 
under  the  hood  at  a  distance  from  any  woodwork,  and  start 
the  reaction  by  igniting  the  end  of  the  magnesium  ribbon 
with  a  Bunsen  flame.  It  is  advisable  for  the  operator  to 
wear  colored  glasses  while  watching  the  reaction,  and  to 
keep  at  a  little  distance  to  be  out  of  the  way  of  flying 
sparks.  When  the  crucible  has  cooled,  break  it  and  sepa- 
rate the  regulus  of  metallic  chromium  from  the  slag  of 
fused  aluminum  oxide. 


38.     MANGANESE    CHLORIDE    FROM    WASTE 
MANGANESE    LIQUORS 

The  waste  liquors  left  after  the  generation  of  chlorine 
from  manganese  dioxide  and  hydrochloric  acid  contain 
principally  manganous  chloride.  Besides  this,  however, 
there  is  always  some  free  acid  and  almost  always  a  con- 
siderable amount  of  ferric  chloride  present.  The  greater 
part  of  the  free  acid  can  be  removed  by  evaporating  the 


MANGANESE    CHLORIDE 

solution  until  a  pasty  mass  is  left  which  will  solidify  on 
cooling.  The  iron  can  be  removed  from  the  solution  of 
this  residue  in  virtue  of  the  ease  with  which  ferric  salts 
hydrolyze.  The  nearly  neutral  solution  is  treated  with 
suspended  manganous  carbonate  (obtained  by  treating  a 
part  of  the  solution  itself  with  a  soluble  carbonate).  Ferric 
chloride  hydrolyzes  according  to  the  reversible  reaction, 
FeCl3  +  3H2O  ^  3HC1  +  Fe(OH)8.  In  the -presence  of 
manganous  carbonate  the  small  amount  of  free  acid  thus 
formed  is  continuously  used  up  according  to  the  reaction, 
MnCO3  +  2HC1  -*  MnCl2  +  H2O  +  CO2.  Thus  the  re- 
action of  hydrolysis  is  enabled  to  run  to  completion.  The 
remaining  solution,  which  is  almost  absolutely  neutral  and 
entirely  free  from  iron  salts,  yields  crystallized  manganous 
chloride,  MnCl2.4H2O,  upon  evaporation. 
*  Procedure.  —  Boil  500  cc.  of  waste  manganese  liquor  in 
a  6-inch  evaporating  dish  under  the  hood  until  the  residue 
becomes  pasty.  After  a  scum  begins  to  form  on  the  surface 
of  the  liquid,  there  is  danger  of  spattering  and  the  mixture 
should  be  stirred  with  a  glass  rod  until  it  becomes  semi- 
solid.  Heat  the  residue  to  boiling  with  1,000  cc.  of  water- 
without  filtering  the  solution  obtained,  take  one-tenth  of  it, 
dilute  this  portion  to  1,000  cc.,  and  add  a  solution  of  sodium 
carbonate  to  it  until  all  of  the  manganese  is  precipitated  as 
carbonate  (test  for  complete  precipitation).  Transfer  the  pre- 
cipitate to  a  tall,  common  bottle  and  wash  it  by  decantation 
at  least  four  times.  Add  the  slime  of  manganous  carbonate 
to  the  remaining  nine-tenths  of  the  manganous  chloride 
solution,  and  boil  the  mixture  in  a  casserole  until  a  few 
drops  of  the  filtered  liquid  give  no  red  color  when  tested 
with  potassium  sulphocyanate.  Filter  the  solution  and 
evaporate  it  in  an  8-inch  dish  until  a  crystal  scum  forms 
on  blowing  across  the  surface.  Then  allow  the  solution  to 
cool  slowly  and  crystallize,  leaving  it  for  at  least  12  hours 


I3O  HEAVY     METALS 

uncovered  in  a  place  protected  from  dust  Collect  the 
crystals  and  evaporate  the  mother  liquor  to  obtain  further 
crops  of  crystals  until  practically  all  of  the  salt  has  crystal- 
lized.1 Spread  the  light  pink  crystals  on  an  unglazed  plate 
to  dry. 

Questions 

1.  Explain    the   purpose   of    the    test   with   potassium 
sulphocyanate. 

2.  Explain   the   action    of   manganese    dioxide   in    the 
generation    of    chlorine    gas    from    hydrochloric    acid.      In 
what  state  of  oxidation  does  manganese  exist  in  the  salt 
manganous  chloride  ? 

3.  If  iron  were  in  the  ferrous  condition,  it  would  not 
be   removed   from   the   solution   by   the    above   procedure. 
Explain  why  iron   is   necessarily  in   the  ferric  condition  in 
the  liquors  used. 

4.  Experiment.  —  Dissolve  a  small  grain  of  manganous 
chloride  in  a  half  test  tube  of  water.     Test  the  solution  with 
hydrogen  sulphide ;  then  add  a  few  drops  of  ammonia,  and 
if  necessary  pass  in  a  little  more  hydrogen  sulphide.     Then 
add  acetic  acid  (a  weak  acid)  until  the  solution  is  again 
faintly  acid.     Does  the  manganous  sulphide  dissolve  ?    Com- 
pare  the    solubility   of    manganous    sulphide   with   that   of 
copper  sulphide ;  of  zinc  sulphide. 

5.  In   testing   for   the    complete   precipitation  of    iron 
from  the  manganese  chloride  solution,  what  would  be  the 
effect  observed  on  adding  ammonium  sulphide  (a)  before, 
and  (b)  after,  all  the  iron  has  been  precipitated  ? 

JThe  crystals  of  manganous  chloride  are  deliquescent  when  the 
temperature  is  low  and  the  atmosphere  charged  with  moisture.  If  the 
product  cannot  be  obtained  satisfactorily  by  the  above  directions, 
carry  out  the  crystallization  and  drying  in  a  place  at  a  slightly  elevated 
temperature,  25-30°;  or  cool  the  saturated  hot  solution  rapidly  by 
stirring  or  shaking,  and  dry  the  crystal  meal  so  obtained  by  rinsing  it 
with  alcohol  and  then  letting  the  latter  evaporate  rapidly. 


POTASSIUM    PERMANGANATE  13! 

6.  Explain  how  facts  involved  in  the  foregoing  prepa- 
ration show  that  Mn(OH)2  is  more  strongly  basic  than 
Fe(OH)3. 

39.     POTASSIUM    PERMANGANATE    FROM 
MANGANESE    DIOXIDE 

Although  manganese  dioxide  is  a  powerful  oxidizing 
agent,  it  is  nevertheless  capable  of  being  itself  oxidized 
when  it  is  fused  with  a  basic  flux.  The  trioxide  of  manga- 
nese is  acidic  in  nature  and  combines  with  the  base  to  form 
a  salt.  Thus  it  is  evident  that  the  presence  of  a  base  favors 
the  oxidation. 

The  dioxide  of  manganese  is  neither  strongly  basic  nor 
acidic  in  nature  and  shows  no  marked  tendency  to  form  salts. 
The  monoxide  is  distinctly  basic  and  the  trioxide  is  distinctly 
acidic,  so  that  the  former  forms  salts  with  acids  and  the 
latter  with  bases.  It  follows,  therefore,  that  in  the  presence 
of  acids  the  dioxide  has  a  tendency  to  produce  salts  of  man- 
ganous  oxide  whereby  an  atom  of  oxygen  is  set  free  (see 
No.  38,  Manganous  Chloride),  and  that  in  the  presence  of 
bases  manganese  dioxide  has  a  tendency  to  take  on  another 
atom  of  oxygen  in  order  to  produce  a  salt  of  the  trioxide. 

Thus  when  manganese  dioxide  is  fused  with  potassium 
hydroxide  and  an  oxidizing  agent,  the  salt  potassium  man- 
ganate  is  formed.  This  salt  is  soluble  in  water  and  is  fairly 
stable  so  long  as  a  considerable  excess  of  potassium  hydrox- 
ide is  present ;  but  in  presence  of  an  acid  —  even  so  weak 
a  one  as  carbonic  acid  —  the  manganate  decomposes  spon- 
taneously, two-thirds  being  oxidized  to  permanganate  at 
the  expense  of  the  other  one-third,  which  is  reduced  again 
to  manganese  dioxide : 

3H2MnO4  -*  2HMnO4  +  MnO2  -f-  2H2O. 
The   permanganate   (or    permanganic   acid)  corresponds   to 


132  HEAVY     METALC 

the  heptoxide  of  manganese,  Mn2O7,  which  is  the  most 
strongly  acid-forming  of  the  oxides  of  manganese.  Perman- 
ganic acid  is  a  strong  and  very  soluble  acid,  it  being  of 
approximately  the  same  acid  strength  as  nitric  or  hydro- 
chloric acids.  It  is  in  addition  a  very  powerful  oxidizing 
agent. 

Procedure. —  Place  50  grams  of  potassium  hydroxide 
and  25  grams  of  potassium  chlorate  in  an  8  cm.  sheet  iron 
crucible.  Heat  the  mixture  carefully  until  it  is  just  melted. 
Meantime  grind  50  grams  of  pyrolusite  to  as  fine  a  powder 
as  possible  (the  finer  it  is  ground,  the  more  successful  the 
preparation).  Remove  the  flame  from  under  the  crucible 
and  add  the  pyrolusite,  a  little  at  a  time,  stirring  vigorously 
with  an  iron  spatula  (an  old  file  with  a  wooden  handle)  all 
the  while.1  After  all  is  added,  place  a  small  flame  below 
the  crucible,  and  keep  stirring  the  charge.  Gradually  in- 
crease the  strength  of  the  flame,  and  stir  continuously  until 
the  mass  stiffens  completely.  Then  cover  the  crucible  and 
heat  it  5  minutes  longer  at  a  dull  red  heat.  When  the  mass 
has  cooled,  place  crucible  and  all  in  i  liter  of  water  in  an 
8-inch  porcelain  dish.  After  the  solid-  has  entirely  disin- 
tegrated, remove  the  crucible  and  rinse  it  off  with  a  little 
water  from  the  wash  bottle.  Boil  the  solution  in  the  dish, 
and  at  the  same  time  pass  in  carbon  dioxide  generated  from 
marble  and  hydrochloric  acid  until  the  green  color  of  the 
manganate  has  entirely  changed  to  the  violet-red  color  of 
the  permanganate.  Test  the  color  by  touching  a  drop  of 
the  solution  on  a  stirring  rod  to  a  piece  of  filter  paper. 
If  the  spot  is  violet  with  no  trace  of  green  and  only  a  fleck 
of  brown  manganese  dioxide  in  the  center,  the  change  to 

1  Since  the  charge  in  the  crucible  effervesces  and  spatters  particles 
of  melted  salt,  great  care  should  be  taken  to  keep  the  eyes  at  a  safe 
distance.  The  hand  holding  the  stirrer  should  be  protected  with  a 
thick  glove  or  with  a  towel,  and  with  the  other  hand  the  crucible 
should  be  held  firmly  by  means  of  long  iron  tongs. 


POTASSIUM    PERMANGANATE  133 

permanganate  is  complete.  Remove  the  lamp ;  let  the  sludge 
settle  in  the  dish  for  5  minutes ;  then  pour  the  solution 
through  an  asbestos  filter  (see  Note  4  (//)  on  page  10),  being 
careful  to  avoid  stirring  up  the  sludge  until  the  very  last, 
since  the  slimy  precipitate  of  manganese  dioxide  would  so 
clog  the  filter  as  to  nearly  stop  the  flow.  Lastly,  with  the 
aid  of  a  jet  of  water  from  the  wash  bottle,  transfer  all  the 
sludge  to  the  filter  and  drain  it  free  from  liquid.  Evaporate 
the  solution  in  a  clean  dish  to  a  volume  of  300  cc.  Let  it 
settle  a  moment  and  filter  it  through  asbestos  as  before. 
Pour  the  filtrate  into  a  6-inch  dish,  and  allow  it  to  cool 
slowly  in  a  place  protected  from  the  dust.  When  cold,  col- 
lect the  crystals  of  potassium  permanganate  on  a  perforated 
plate  placed  loosely  in  a  filter  funnel.  Evaporate  the  mother 
liquor  to  100  cc.,  filter  it  through  asbestos,  and  obtain  a 
second  crop  of  crystals.  Discard  the  remaining  liquid,  since 
it  cannot  contain  more  than  about  6  grams  of  potassium 
permanganate  and  to  evaporate  it  further  would  cause  po- 
tassium chloride  also  to  crystallize  out.  Weigh  all  the  crys- 
tals, dissolve  them  in  eight  times  their  weight  of  water  (to 
give  a  saturated  solution  at  about  40°),  filter  the  solution 
through  asbestos  at  near  the  boiling  temperature,  and  let 
it  cool  slowly  and  crystallize  in  a  small  porcelain  dish  cov- 
ered with  a  watch  glass.  Recover  another  crop  of  crystals 
in  the  same  way  from  the  mother  liquor,  after  evaporating 
it  to  a  volume  of  60  cc.  Allow  the  crystals  to  dry  on  a 
clean  unglazed  plate. 

Questions 

1.  Name  and  give  the  symbols  of   all  the  oxides  of 
manganese. 

2.  From  which  oxide  is  K2MnO4  derived  ?     KMnO4  ? 

3.  Write  the  reactions  involved  in  the  above  preparation. 

4.  How  could  KMnO4  be  converted  back  into  K2MnO4? 
Reaction  ? 


134  HEAVY    METALS 

5.  Does  it  frequently  happen  that,  with  an  element 
which  can  exist  in  several  states  of  oxidation,  a  compound 
derived  from  one  oxide  is  stable  in  an  alka'line  solution  but 
unstable  in  an  acid  solution,  while  in  the  latter  solution 
the  compound  derived  from  another  oxide  is  the  stable  one  ? 
What  other  preparation  besides  the  present  one  illustrates 
this  point? 


40.     MANGANESE    METAL    BY    THE    GOLDSCHMIDT 
PROCESS 

The  principle  of  the  production  of  manganese  by  this 
process  is  exactly  the  same  as  that  of  the  production  of 
chromium  in  Exercise  33.  On  account  of  the  violence  of 
the  reaction  between  the  oxide  of  manganese  and  aluminum 
it  is  not  advisable  to  ignite  the  whole  charge  at  once  in  the 
crucible ;  yet  on  account  of  the  high  melting  point  of  man- 
ganese a  considerable  quantity  of  charge  must  be  used  in 
order  to  produce  heat  enough  to  obtain  the  metal  melted 
together  in  a  uniform  lump,  instead  of  distributed  in  small 
globules  throughout  the  mass  of  the  slag.  Before  mixing 
up  the  charge,  the  pyrolusite  which  is  used  must  be  first 
heated  by  itself  in  order  to  drive  off  any  water  which  it  may 
contain  and  to  convert  it  to  the  lower  oxide,  Mn2O3. 

Procedure.  —  Place  i  kilogram  of  finely  powdered  pyro- 
lusite in  a  Hessian  crucible  and  heat  to  a  bright  heat  in  a 
gas  furnace.  To  prepare  the  charge,  mix  750  grams  of 
this  material,  when  it  is  cooled  sufficiently,  with  250 
grams  of  granulated  aluminum.  Heat  the  empty  crucible 
again  in  the  furnace,  and  while  still  hot  place  it  in  a  pail 
of  sand,  as  in  the  preparation  of  chromium.  Place  about 
20  grams  of  the  charge  in  the  bottom  of  the  hot  crucible. 
Put  on  colored  glasses  and  a  heavy  glove ;  start  the  reac- 
tion with  a  little  fuse  powder  and  a  magnesium  ribbon  (see 


FERROUS    AMMONIUM    SULPHATE  135 

Chromium),  and  then  add  fresh  portions  of  the  charge 
rapidly  but  without  allowing  the  reaction  to  become  too 
violent.  When  the  crucible  has  cooled,  break  it,  and  sepa- 
rate the  regulus  of  metallic  manganese  from  the  slag  of 
fused  aluminum  oxide. 

Questions 

1.  If   pyrolusite   containing   water   were    used  without 
previous   heating,  what   disadvantage  would   result   during 
the  process? 

2.  What  economy  of  materials  is  effected  by  converting 
the  manganese  dioxide  into  the  lower  oxide  ? 


41-     FERROUS    AMMONIUM    SULPHATE 
FeSO4.(NH4)2SO4.6H20 

Corresponding  to  the  two  most  important  oxides  of  iron, 
FeO  and  Fe2O3,  the  two  sulphates,  FeSO4  and  Fe2(SO4)8, 
can  be  prepared.  By  dissolving  iron  in  sulphuric  acid  a 
solution  of  ferrous  sulphate  is  obtained.  This,  however, 
is  readily  oxidizable,  slowly  even  by  the  oxygen  of  the  air, 
to  the  higher  sulphate,  and  ferrous  sulphate  can  only  be 
preserved  free  from  ferric  salt  when  all  oxygen  is  excluded, 
or  when  it  is  kept  in  contact  with  an  excess  of  metallic 
iron  in  an  acidified  solution.  Dry  crystallized  ferrous  sul- 
phate or  green  vitriol,  FeSO^yHoO,  can  be  preserved  fairly 
well  without  becoming  oxidized  ;  but  the  double  ferrous  and 
ammonium  sulphate  is  not  only  more  easily  prepared  on 
account  of  the  readiness  with  which  it  crystallizes,  but  it 
is  also  much  less  easily  oxidized  by  contact  with  the  air. 

Procedure. —  Prepare  crystallized  ferrous  ammonium  sul- 
phate from  equimolal  quantities  of  crystallized  ferrous  sulphate 
and  ammonium  sulphate,  using  70  grams  of  the  former  and 


136 


HEAVY     METALS 


33  grams  of  the  latter.  In  crystallizing  the  product,  observe 
the  solubilities  given  in  the  following  table,  and  make  use 
of  suggestions  given  in  Note  8,  page  15  and  under  Alum, 
page  54: 

A  saturated  solution  contains  for  each  100  grams  of  water  the  given 
number  of  grams  of  the  anhydrous  salt. 


0° 

10° 

20° 

30° 

40° 

50° 

70° 

90° 

FeSO4     

16 

26 

33 

44 

48 

56 

43 

(NtDoSO.. 

71 

73 

75 

78 

81 

84 

92 

99 

FeSO4  (NH4)2SO4  

12 

17 

22 

28 

33 

40 

52 

Questions 

i.  Dissolve  a  little  of  the  preparation  in  water  and  test 
it  with  potassium  ferrocyanide.  If  the  precipitate  is  white 
or  only  a  pale  blue,  of  what  does  it  consist  ?  If  it  is  deep 
blue,  what  is  shown? 


42.     FERRIC    AMMONIUM    ALUM 

In  this  preparation  ferrous  sulphate  is  converted  into 
ferric  sulphate  under  the  oxidizing  action  of  nitric  acid  in 
the  presence  of  the  amount  of  sulphuric  acid  theoretically 
necessary  to  form  this  salt.  By  the  addition  of  ammo- 
nium sulphate  the  double  salt,  ferric  ammonium  sulphate, 
Fe2(SO4)8,(NH4)2"SO4.24H2O,  crystallizes,  this  being  one  of 
the  isomorphous  series  of  alums  (see  Alum). 

At  25°,  100  grams  of  water  dissolve  44  grams  of  the 
anhydrous  or  124  grams  of  the  hydrated  ferric  ammonium 
sulphate. 

Procedure.  —  Heat  together  100  grams  of  crystallized 
ferrous  sulphate,  100  cc.  of  water,  and  12  cc.  of  concen- 
trated sulphuric  acid  until  the  salt  is  dissolved.  While  the 


FERRIC    ALUM  137 

solution  is  boiling,  add  concentrated  nitric  acid,  a  little  at 
a  time,  until  the  iron  is  completely  oxidized  to  ferric  sul- 
phate and  a  few  drops  of  the  solution  diluted  with  a  few 
cubic  centimeters  of  water  give  no  blue  precipitate  with 
potassium  ferricyanide.  Evaporate  the  solution  until  it  is 
thick  and  sticky  and  most  of  the  excess  of  nitric  acid  has 
been  driven  off.  Dissolve  this  in  water,  making  up  to  a 
volume  of  125  cc. ;  heat  to  boiling  and  add  25  grams  of 
ammonium  sulphate  dissolved  in  100  cc.  of  hot  water. 
Allow  the  solution  to  cool  slowly  and  crystallize.  Collect 
the  crystals  in  a  funnel ;  wash  with  a  very  little  water  and 
allow  to  dry  on  an  unglazed  plate.  Obtain  a  second  crop 
of  crystals  from  the  mother  liquor. 

Questions 

1.  Write   the   reaction    involved    in   the    oxidation   of 
ferrous  sulphate  as  carried  out  in  this  preparation. 

If  an  unacidified  solution  of  ferrous  sulphate  is  oxidized 
by  the  oxygen  of  the  air,  what  products  are  formed  ? 

2.  Write  the  reaction  involved  in  the  test  for  ferrous 
salt  with  potassium  ferricyanide. 

3.  Experiment.  —  Prepare   a   solution  of  a  ferrous  salt 
by  dissolving   2   grams   of  ferrous   ammonium   sulphate  in 
20  cc.  of  water,  adding  a  little  dilute  sulphuric  acid  and  a 
piece  of  iron  wire.     Test  both  this  solution  and  a  solution 
of  a  ferric  salt  (nitrate  or  chloride)  with  potassium  ferrocya- 
nide,  potassium  ferricyanide,  and  potassium  sulphocyanate. 
Tabulate  the  results.     These  constitute  the  standard  tests 
for  ferrous  and  ferric  salts. 


138  HEAVY    METALS 

GENERAL    QUESTIONS.     VI 

HEAVY    METALS    OF   THE   SIXTH,    SEVENTH,    AND    EIGHTH    GROUPS 
OF   THE    PERIODIC    SYSTEM 

Experiments 

(The  results  observed  are  to  be  recorded  in  the  laboratory  note- 
book at  the  time  the  experiments  are  performed.) 

1.  Test   the    stability  of   nickel   carbonate  by  heating 
i  gram  of  it  gently  in  a  test  tube  while   shaking   it   in  a 
Bunsen  flame.     Test  the  gas   evolved  for  carbon  dioxide; 
and  compare  the  action  of  the  remaining  solid,  when  treated 
with  hydrochloric  acid,  with  that  of  the  original  carbonate. 
The  carbonates  of   divalent    iron,   cobalt,   manganese,   and 
chromium   are   all   of    approximately   the    same    degree   of 
stability  as  nickel  carbonate,  so  that  this  one  experiment 
may  be  taken  as  typical  of  this  class  of  carbonates. 

2.  To  Show  Whether  the  Carbonate  of  a  Trivalent  Metal 
Can  Exist.  —  Dissolve  2  grams  of  ferric  alum  in   i  o  cc.  of 
water  (this  gives  a  trivalent  iron  salt  in  a  solution  that  con- 
tains no  free  acid).     Add  a  10  per  cent,  sodium  carbonate 
solution  slowly  until  no  more  action  takes  place.     What  is 
the    gas  evolved  ?     What    is    the    precipitate  ?     In  this  ex- 
periment the  ions  Fe+  +  + and  CO3~ ~  are  brought  together; 
the  other  ions,  Na+  and  SO4~  ~,  could  not  react  together  to 
give   any  visible  effect.     If,   therefore,   ferric  carbonate    is 
stable  in  contact  with  water,  it  will  either  form  a  precipitate 
if  it  is  insoluble,  or  if  it  is  soluble   it  will  simply  stay  in 
solution  and  no  effect  will  be  observable.     The  gas  given 
off  shows  that  the  carbonate  is  unstable.    Write  the  equation 
for  the  reaction.     None  of  the  carbonates  of  the  metals  of 
this  group,  when  they  are  in  the  trivalent  condition,  are  any 
more   stable  than  ferric   carbonate.     A  salt  of   chromium, 
such  as  chromic  alum,  might  be  used  instead  of  ferric  alum 
in  the  above  experiment. 


GENERAL    QUESTIONS    VI  139 

3.  Oxidation  of  a  Divalent   Oxide.  —  Heat  \  gram  of 
cobalt   carbonate    in   an  open   porcelain    dish,   holding   the 
dish  with  tongs  and  keeping  it  rotating  in  the  flame,  but 
not  allowing  the  porcelain  to  even  approach   a  visible  red 
heat.     Heat    until   the  color  of   the  cobalt   carbonate    has 
completely   changed.      Like    nickel   carbonate,    cobalt    car- 
bonate is  decomposed  by  heat  into  cobaltous  oxide,  CoO, 
and    carbon    dioxide.      If    the    cobaltous    oxide    is    readily 
oxidized   by   the    oxygen    of   the    air,    it    may    at    once   be 
changed    into   Co2O3  or   Co3O4.     To  test  for  this,   treat   a 
little  of  the  product  with  hydrochloric  acid  in  a  test  tube. 
From  CoO  the  salt  CoCl2  would  be  obtained.     From  Co2O8 
or   Co3O4  the   same   salt,   CoCl2,  would  be   obtained  —  not 
CoCl3  —  and  chlorine  would  be  liberated  (action  similar  to 
that  of  MnO2  with  hydrochloric  acid).     Test  for  the  forma- 
tion of  chlorine  by  means  of  the  odor  or  by  using  iodide 
starch  paper.     NOTE  :  In   Experiment  i   a  higher  oxide  of 
nickel  was  probably  formed  in  the  same  manner,  although 
to  a  considerably  less  extent. 

4.  Acidify  solutions   of    potassium    permanganate    and 
potassium  bichromate  each  with  sulphuric  acid,  warm  them 
and  treat  them  with  sulphur  dioxide  (sulphurous  acid),  and 
note  any  change  in  color.     The  change  is  due  to  the  reduc- 
tion of  the  given  salts  which  are  derived   from  the  oxides 
Mn2O7   and   CrO3,   respectively,  to   salts   derived    from  the 
oxides  MnO  and  Cr2O3,  respectively. 

Questions 

In  the  group  of  elements  discussed  under  General  Questions,  I, 
changes  of  valence  do  not  occur,  but  the  metals  of  the  alkali  and  alka- 
line earth  families  show  the  same  valence  in  all  their  compounds. 
Proceeding  in  the  order  in  which  the  elements  have  been  taken  up 
in  this  book,  a  constantly  increasing  tendency  is  observed  for  the  ele- 
ments to  display  different  valences,  until  in  the  group  under  consid- 
eration the  most  important  chemical  characteristics  of  the  elements 


I4O  HEAVY     METALS 

are  connected  with  their  changes  from  one  state  of  valency  to  another. 
NOTE  :  The  terms  state  of  valency  and  state  of  oxidation  can  in  most 
cases  be  used  interchangeably. 

1.  In    which   groups   of    the    periodic    system    do   the 
elements    chromium,    manganese,    iron,    nickel,    and    cobalt 
fall  ?     What  is  peculiar  about  the  position  of  the  last  three  ? 
What  other  metals  belong  to  the  same  family  as  chromium  ? 
In  what  relation  do  they  stand  to   sulphur,  selenium,  and 
tellurium  ?     In  what  relation  does  manganese  stand  to  the 
halogens  ?     What  other  elements  occur  in  the  eighth  group 
in  triads  similar  to  iron,  nickel,  and  cobalt  ? 

2.  How  do  the   monoxides  of   chromium,   manganese, 
iron,  cobalt,  and  nickel  compare  in  basic  strength  with  the 
oxides  of  copper  and  zinc  and  with  the  oxides  of  the  alkali 
and  alkaline  earth  metals  ?     How  do  the  sesquioxides,  R2O3, 
compare  with  the  monoxides  of  this  group  as  regards  basic 
strength  (see  Experiment  2)  ? 

What  is  true  as  regards  the  base-  or  acid-forming  prop- 
erties of  the  oxides  higher  than  the  sesquioxides,  e.g.,  of 
CrO3,  MnO3,  Mn2O7  ? 

3.  Give  the  symbols  and  names  of  salts  derived  from 
each  of  the  three  oxides  of  chromium,  CrO,  Cr.2O3,  CrO3. 
In  which  of  its  compounds  does  chromium  most  resemble 
sulphur  ?   iron   and   aluminum  ?   nickel,   cobalt,   copper,  and 
zinc  ? 

4.  Give  the  symbols  and  names  of  salts  derived  from 
each   of   the  oxides   of   manganese,    MnO,   Mn2O3,   MnO2, 
MnO3,  Mn2O7.     In  which  of  its  compounds  does  manganese 
most  resemble  chlorine  ?  aluminum  ?  cobalt,  nickel,  copper, 
and  zinc  ?  sulphur  ?  lead  in  the  dioxide  ? 

5.  Formulate  the  reaction  between  sulphurous  acid  and 
potassium  permanganate  in  acid  solution  (see  Experiment  4). 
By  means  of  partial  equations  resolve  the  compounds  into 
their   constituent   oxides ;    show   the   simple   oxidation  and 


GENERAL     QUESTIONS    VI  14! 

reduction,  and  then  show  how  the  new  oxides  combine  to 
give  the  salts  which  actually  result.  Then  add  the  separate 
equations  to  give  the  total  equation  for  the  complete  reaction. 
6.  In  the  same  manner  formulate  the  reaction  between 
sulphurous  acid  and  potassium  bichromate. 


CHAPTER   VII 

NON-METALLIC    ELEMENTS   OF   THE    SIXTH    AND 
SEVENTH  GROUPS  OF  THE  PERIODIC  SYSTEM 

The  elements  which  are  distinctly  and  invariably  non- 
metallic  in  character  are  boron  in  the  third  group,  carbon 
and  silicon  in  the  fourth  group,  nitrogen  and  phosphorus 
in  the  fifth  group,  oxygen,  sulphur,  selenium,  and  tellurium 
in  the  sixth  group,  and  fluorine,  chlorine,  bromine,  and 
iodine  in  the  seventh  group.  It  has  been  assumed  that 
these  elements,  or  at  least  the  most  important  of  them,  were 
studied  before  entering  on  the  course  of  study  outlined  in 
this  book.  Indeed  no  knowledge  of  the  chemistry  of  the 
metallic  elements  would  be  possible  without  a  certain  knowl- 
edge of  the  non-metallic  elements  with  which  they  form 
compounds. 

By  turning  to  the  table  of  the  periodic  arrangement  of 
the  elements,  it  is  at  once  seen  that  the  non-metals  do  not 
occur  at  all  in  the  first  and  second  groups ;  that  they  occur 
only  at  the  top  in  the  third,  fourth,  and  fifth  groups ;  and 
that  in  the  sixth  and  seventh  groups  they  comprise  all  the 
members  of  the  B  families.  It  is  true  in  these  families,  as 
might  be  expected  by  recalling  characteristics  of  preceding 
groups,  that  the  strength  of  the  non-metallic  character  grows 
weaker  and  that  the  approach  towards  metallic  character 
grows  more  evident,  as  the  atomic  weight  increases ;  indeed 
it  is  probable  that  if  the  elements  which  should  fit  into  the 
places  below  tellurium  and  iodine,  respectively,  are  ever 
found,  they  will  display  quite  as  marked  metallic  properties 
as  non-metallic. 

H3 


144  NON-METALLIC    ELEMENTS 

The  characteristic  valences  of  the  sixth  and  seventh 
groups  are  VI  and  VII,  respectively,  and  the  corresponding 
oxides  are  EO8  and  E2O7.  In  these  oxides  and  in  the 
compounds  derived  from  them,  there  is  little  dissimilarity 
between  the  A  and  B  families.  Thus  perchlorates  and 
permanganates  are  in  every  way  analogous  to  each  other, 
as  are  also  sulphates  and  chromates.  In  any  lower  state 
of  valence,  the  elements  of  the  B  families  are  entirely  dif- 
ferent from  those  of  the  A  families,  and  the  most  striking 
non-metallic  properties  of  the  former  are  exhibited  in  their 
ability  to  combine  directly  with  metallic  elements,  forming 
oxides,  sulphides,  chlorides,  bromides,  etc.  It  will  be  noticed 
that  the  negative  valence,  that  is,  the  valence  exhibited 
towards  hydrogen  or  a  metal,  is  I  for  the  non-metals  of 
the  seventh  group  and  II  for  the  sixth  group. 


43.     POTASSIUM    IODIDE 

Of  the  two  most  obvious  possibilities  for  making  this 
salt,  the  direct  synthesis  from  the  elements  is  quite  imprac- 
ticable, both  because  potassium  metal  is  expensive  and 
because  the  action  would  be  violent  and  difficult  to  control. 
The  method  of  neutralizing  hydriodic  acid  with  potassium 
hydroxide  presents  no  chemical  difficulties,  but  the  materials 
are  more  costly  than  the  iodine  and  potassium  carbonate 
which  are  employed  according  to  the  following  procedure 
(compare  with  the  preparation  of  ammonium  bromide, 
No.  5). 

Iodine,  when  brought  together  in  the  presence  of  water 
with  an  excess  of  iron,  reacts  to  form  soluble  ferrous  iodide. 
By  treating  this  solution  with  potassium  carbonate  a  metath- 
esis takes  place,  yielding  insoluble  ferrous  carbonate  and 
soluble  potassium  iodide.  The  desired  product  should  then 
be  obtained  by  filtration,  except  that  the  ferrous  carbonate 
forms  a  slimy,  bulky  precipitate  which  clogs  the  filter;  but 
this  difficulty  may  be  overcome  if  the  ferrous  salt  is  partially 
oxidized  by  adding  more  iodine  before  throwing  out  the 
iron,  and  the  mixture  of  ferrous  carbonate  and  ferric  hydrox- 
ide that  is  obtained  in  this  way  will  filter  very  readily. 

Procedure.  —  Place  7  or  8  grams  of  iron  filings  and 
«jo  cc.  of  water  in  an  Erlenmeyer  flask  and  add  25  grams 
of  iodine,  a  small  portion  at  a  time,  while  shaking  continu- 
ously. When  all  is  added  warm  the  mixture  somewhat  until 
ail  of  the  iodine  has  combined,  and  the  dark  brown  color 
of  free  iodine  gives  place  to  the  yellow  of  ferrous  iodide. 
Filter  off  the  excess  of  iron  and  add  5  grams  more  of  iodine 
to  the  solution.  Warm  the  solution  until  the  iodine  is  dis- 
solved, and  then  pour  it  into  a  boiling  solution  of  17  grams 
of  potassium  carbonate  in  50  cc.  of  water  in  a  good-sized 

MS 


146  NON-METALLIC    ELEMENTS 

flask.  Warm  the  solution,  which  is  at  first  very  thick  with 
the-  gelatinous  precipitate,  until  the  latter  becomes  more 
compact  and  the  mixture  thus  becomes  more  fluid.  Filter 
a  little  of  the  liquid ;  it  should  be  colorless  and  should  con- 
tain no  iron;  otherwise  a  little  more  potassium  carbonate 
should  be  added.  Filter  the  whole  solution  and  wash  the 
precipitate  with  hot  water  to  save  all  the  soluble  salt. 
Evaporate  the  filtrate  to  a  small  volume  in  a  porcelain  dish, 
filter  again  if  necessary,  and  evaporate  further  until  crystals 
begin  to  form.  Then  leave  the  solution  to  evaporate  spon- 
taneously in  a  place  protected  from  dust  (best  in  a  somewhat 
warm  place).  Collect  the  crystals  and  recover  another  crop 
from  the  mother  liquor. 

Questions 

1.  What  proportion  of  the  ferrous  compound  is  oxidized 
to  ferric  by  the  addition  of  the  5  grams  of  iodine  to  the 
filtered  solution  in  the  foregoing  procedure  ? 

2.  On  treating  the  mixture  of  ferrous  iodide  and  iodine 
with  potassium  carbonate,  state  reasons  why  ferrous  carbon- 
ate rather  than  ferrous  hydroxide  and  why  ferric  hydroxide 
rather  than  ferric  carbonate  should  precipitate. 

3.  Starting  with  iodine  and  potassium  hydroxide,  devise 
a  process  for  preparing  potassium  iodide  without  the  use  of 
iron  or  similar  metal.     Compare  No.  5,  Ammonium  Bromide, 
and  No.  46,  Potassium  Bromide  and  Potassium  Bromate. 


44.     HYDRIODIC   ACID 

The  direct  synthesis  of  hydrogen  iodide  from  the  ele- 
ments is  impracticable,  because  the  chemical  affinity  between 
hydrogen  and  iodine  is  so  small  that  at  any  temperature 
sufficiently  elevated  to  make  them  react  at  all,  they  would 
combine  only  very  incompletely.  By  the  interaction  of  an 


HYDRIODIC    ACID  147 

iodide  with  a  non-volatile  acid,  such  as  sulphuric  acid, 
hydrogen  iodide  gas  could,  it  is  true,  be  formed : 

KI  -f  H2S04  -  KHS04  +  HI ; 

but  the  hydrogen  iodide  so  formed  acts  as  a  reducing  agent 
upon  the  sulphuric  acid,  whereby  it  is  itself  oxidized  to  free 
iodine  and  water. 

Hydrogen  iodide,  then,  cannot  be  satisfactorily  prepared 
by  the  direct  union  of  hydrogen  and  iodine  nor  by  the  me- 
tathesis of  an  iodide  with  the  non-volatile  sulphuric  acid. 
The  most  convenient  method  of  preparing  it  is  by  means 
of  the  action  of  hydrogen  sulphide  with  iodine  in  aqueous 
solution : 

H2S+  I2  ->  2HI  -f  S. 

The  affinity  between  hydrogen  and  sulphur  is  very  small, 
as  is  also  that  between  hydrogen  and  iodine ;  but  this  fact, 
combined  with  the  fact  that  free  iodine  is  appreciably  solu- 
ble while  free  sulphur  is  exceedingly  insoluble  in  water, 
makes  it  possible  for  the  reaction  to  run  to  completion. 

The  rather  dilute  solution  of  hydriodic  acid  which  is 
obtained  in  this  manner  can  be  concentrated  by  distillation ; 
at  first  nearly  pure  water  passes  off,  then  the-  quantity  of 
acid  in  the  distillate  increases  until  an  acid  of  specific 
gravity  1.7,  containing  57  per  cent,  of  HI,  comes  over.  At 
this  point  the  distillate  has  the  same  composition  as  the 
residual  liquid,  and  the  remainder  of  the  acid  can  be  dis- 
tilled with  a  constant  composition.  By  this  method  it  is 
not  possible  to  obtain  an  acid  of  higher  concentration,  but 
by  allowing  this  acid  to  absorb  hydrogen  iodide  gas  until  it 
is  saturated  at  o°,  an  acid  of  specific  gravity  2.0,  containing 
90  per  cent,  of  HI,  can  be  obtained.  Hydrogen  iodide  gas 
can  be  prepared  by  the  interaction  of  red  phosphorus  and 
iodine  with  a  little  water. 

Procedure.  —  Grind  30  grams  of  iodine  to  a  fine  powder 


148  NON-METALLIC    ELEMENTS 

and  add  £  gram  of  it  to  100  cc.  of  water  in  a  small  Erlen- 
meyer  flask.  Pass  hydrogen  sulphide  into  the  solution  until 
the  brown  color  of  iodine  has  disappeared.  Add  about 
i  gram  more  of  iodine,  and  again  pass  hydrogen  sulphide 
until  the  iodine  is  used  up.  After  10  grams  of  iodine 
have  reacted  in  this  way,  add  the  remaining  20  grams  and 
allow  the  mixture  to  stand,  with  repeated  shaking,  until  the 
iodine  is  entirely  dissolved  (half  an  hour  or  more).  Then 
pass  hydrogen  sulphide  slowly  until  the  solution  is  decolor- 
ized. Pour  the  solution  into  another  flask,  leaving  the  clotted 
lumps  of  sulphur  behind,  and  rinse  the  first  flask  and  the 
residue  with  a  few  cc.  of  water.  Pass  a  current  of  carbon 
dioxide  through  the  solution  until  the  excess  of  hydrogen 
sulphide  is  entirely  removed ;  then  shake  the  flask  vigor- 
ously to  cause  the  suspended  sulphur  to  clot  together,  and 
filter  the  solution.  In  this  way  a  rather  weak  solution  of 
hydriodic  acid  is  obtained. 

Fit  a  distilling  flask  with  a  thermometer  and  an  inlet 
tube  for  hydrogen,  and  pass  the  side  arm  of  the  flask  into 
a  condenser.  After  introducing  the  hydriodic  acid  solu- 
tion, fill  the  whole  apparatus  with  hydrogen,  and  keep  a 
slow  current  of  this  gas  passing  during  the  distillation. 
On  distilling,  nearly  pure  water  passes  over  at  first  and 
the  thermometer  does  not  register  appreciably  above  100°. 
When  the  thermometer  rises  to  105°  change  the  receiving 
vessel  and  collect  the  distillate  until  the  temperature  has 
risen  to  120°.  Change  the  receiver  again  and  collect  the 
rest  of  the  distillate.  The  temperature  rises  quickly  to 
126°,  and  remains  very  close  to  this  point  until  practically 
all  of  the  acid  has  passed  over.  This  last  fraction  is  the 
desired  concentrated  acid. 

Questions 

i.  Compare  the  stability  of  the  hydrogen  halides  as 
judged  (a)  by  the  heat  produced  when  they  are  formed 


POTASSIUM    CHLORATE  149 

from  the  elements;    and   (b)  by  the   ease  with  which   the 
compounds  are  decomposed. 

2.  Compare   the   action    of    the    hydrogen    halides   as 
reducing  agents.     Does  hydrofluoric  acid  behave   as   a  re- 
ducing  agent    towards   any    substance  ?     What    substances 
are  reduced  by  hydrochloric  acid?     What  substances   can 
be   reduced   by   hydriodic    acid    that   are   not   reduced   by 
hydrochloric  acid  ? 

3.  What  is  the  common  commercial  method  for  prepar- 
ing  hydrochloric   acid  from   sodium  chloride  ?     Why  could 
not  hydrobromic  or  hydriodic  acids  be  prepared  in  a  similar 
manner  ? 

4.  Describe  and  explain  the  method  of  preparing  hydro- 
gen iodide  from  red  phosphorus  and  iodine. 

5.  As  illustrated  in  the  foregoing  preparation  hydriodic 
acid  is  extremely  soluble  in  water.     A  great  deal  of  heat 
also  is  liberated  when  hydrogen  iodide  gas  dissolves  (19,200 
calories  for  each  mol  of  HI).     Compare  this   heat  effect 
with  that  of  some  well-known  chemical  reaction,  for  example, 
the  neutralization  of  a  strong  acid  with  a  strong  base.     In 
the  absence  of  water  the  reaction  2 HI  -|-  S  -*  H2S  -f-  ^2 
takes  place  mainly  in  the  direction  indicated.     What  is  the 
connection  between  the  large  heat  effect  just  mentioned  and 
the  fact  that  the  direction  of  the  reaction  is  reversed  when 
it  takes  place  in  aqueous  solution  ? 


45.     POTASSIUM    CHLORATE 

When  chlorine  dissolves  in  water  it  hydrolyzes  to  some 

extent : 

C12  +  HOH  ^  HC1  +  HOC1  (i) 

The  presence  of  a  base  causes  this  hydrolysis  to  run  to 
completion  because  the  two  acids  produced  by  the  reaction 


ISO  NON-METALLIC    ELEMENTS 

are  immediately  neutralized;  thus  by  passing  chlorine  into 
a  solution  of  sodium  hydroxide  a  mixture  of  chloride  and 
hypochlorite  is  obtained : 

C12  +  2NaOH  -»  NaCl  +  NaOCl  +  H2O.      (2) 

Sodium  hypochlorite  is  fairly  stable  in  a  cold  solution,  but 
in  a  warm  solution  it  is  less  so ;  it  gives  up  its  oxygen,  and 
if  no  more  easily  oxidizable  substance  is  present  it  will  oxi- 
dize either  chloride  or  hypochlorite  ions  to  chlorate  ions : 

3NaOCl  -*  2NaCl  +  NaClO3  (3) 

or     3NaOCl  +  KC1  -»  KC1O3  -f  3NaCl.  (4) 

But  compared  with  a  hypochlorite  (that  is,  the  OCl~-ion),  free 
hypochlorous  acid,  HOC1,  is  a  far  stronger  oxidizing  agent, 
and  therefore  the  formation  of  chlorate  takes  place  more 
readily  when  the  solution  contains  a  trace  of  acid: 

3HC1  +  3NaOCl  -»  3HOC1  +  3NaCl  1 
3HOC1  +  KC1  ->  KC103  +  3HC1        j 

Too  much  acid,  however,  causes  a  reversal  of  reaction  (i), 
HOC1  -f  HC1  -»  H2O  +  C12, 

and  the  action  either  of  chlorine  or  of  hypochlorous  acid  on 
a  strongly  acidified  solution  cannot  produce  any  chlorate. 

One  method  for  preparing  potassium  chlorate  is  to  pass 
chlorine  gas  into  a  hot  concentrated  solution  of  potassium 
hydroxide  until  the  alkali  has  been  entirely  neutralized  and 
a  small  amount  of  free  acid  has  been  formed.  This  point 
is  recognized  by  the  solution  assuming  the  permanent  yellow 
tint  due  to  free  chlorine.  This  method  is  the  one  that  will 
actually  be  used  in  the  preparation  of  potassium  bromate, 
No.  46 ;  but  in  the  present  preparation,  to  avoid  the  use  of 
chlorine  gas,  which  is  very  objectionable  in  the  laboratory, 


POTASSIUM    CHLORATE  151 

bleaching  powder,  or  calcium  hypochlorite,  is  employed  as  the 
oxidizing  agent,  and  potassium  chloride  is  used  in  amount 
only  sufficient  to  supply  the  necessary  potassium  ions. 

Bleaching  powder  is  never  obtained  as  the  pure  sub- 
stance represented  by  the  chemical  symbol ;  it  always  con- 
tains a  considerable  amount  of  unchanged  calcium  hydroxide 
as  well  as  calcium  carbonate.  The  oxidation  to  chlorate 
cannot  be  accomplished  rapidly  so  long  as  there  is  any  con- 
siderable amount  of  calcium  hydroxide  present,  and  this  is 
therefore  neutralized  with  hydrochloric  acid.  After  the  base 
is  neutralized,  any  further  addition  of  acid  begins  to  react 
with  the  calcium  carbonate  and  the  escape  of  carbon  dioxide 
is  observed.  The  carbonic  acid  thus  formed  gives  just 
about  the  right  degree  of  acidity  to  the  solution  to  produce 
the  best  yield  of  chlorate.  If  the  hydrochloric  acid  is  used 
incautiously,  and  more  than  enough  to  react  with  all  the 
carbonate  is  added,  a  corresponding  amount  of  chlorate  will 
be  lost,  for  hydrochloric  acid  reduces  chloric  acid  approxi- 
mately according  to  the  reaction 

HC1  +  KC1O3  ->  KC1  +  HC1O3 
HC1O3  +  HC1  -  H2O  +  Cl  +  C1O2. 

Procedure.  —  Wet  a  mixture  of  175  grams  of  bleaching 
powder  and  20  grams  of  potassium  chloride,  and  rub  it  with 
a  pestle  until  a  smooth,  thin  paste  is  obtained.  Transfer 
this  paste  to  a  tall  bottle  or  jar  and  add  hydrochloric  acid 
(at  the  hood)  as  follows:  Dilute  one  volume  of  acid  of  1.12 
sp.  gr.  with  three  volumes  of  water,  and  add  this  diluted 
acid  very  cautiously  through  a  thistle  tube,  the  stem  of  which 
reaches  to  the  bottom  of  the  liquid.  Rotate  the  bottle  or 
stir  the  contents  vigorously  throughout  the  process  until 
carbon  dioxide  begins  to  be  evolved  freely.  This  point  is 
recognized  when  with  each  fresh  addition  of  \  to  i  cc. 
of  acid  the  gas  that  is  produced  at  the  point  where  the 


152  NON-METALLIC    ELEMENTS 

acid  enters  the  liquid  rises  to  the  top  without  being  absorbed 
and  causes  considerable  frothing.  Probably  about  250  cc. 
of  the  1. 1 2  acid  will  be  required,  but  this  depends  very 
largely  on  tie  quality  of  the  bleaching  powder.  The  reac- 
tion should  now  have  warmed  the  solution  to  about  40°, 
and  the  formation  of  chlorate  is  nearly  completed.  The 
solution  cannot  be  filtered  at  this  point,  as  it  still  contains 
so  much  hypochlorous  acid  that  it  would  disintegrate  the 
filter  paper.  Pour  it,  therefore,  into  an  evaporating  dish 
and  boil  it  (at  the  hood)  until  it  is  concentrated  to  a  volume 
of  about  400  cc.  By  this  time  all  of  the  hypochlorous  acid 
has  either  reacted  to  form  chlorate  or  has  been  otherwise 
decomposed,  and  the  solution  may  be  poured  through  a 
large,  common  filter.  Allow  the  filtrate  to  cool  completely, 
and  collect  the  crystals  of  potassium  chlorate.  Dissolve 
the  crystals  in  hot  water  (see  solubility  table  in  appendix) 
and  recrystallize  twice  or  three  times,  as  may  be  necessary 
to  obtain  the  product  entirely  free  from  chloride  (test  with 
silver  nitrate  solution). 

Questions 

1.  How  is  bleaching  powder  prepared  ?     What  is  its 
formula  and  its  chemical  name? 

2.  Explain   why   the   odor   of    chlorine   becomes   very 
noticeable  when,  according  to  the  foregoing  procedure,  the 
requisite    amount   of    hydrochloric   acid    has    been   added. 
Explain  why  addition  of  any  further  amount  of  acid  'would 
cause  an  appreciable  loss  of  chlorine  and  a  corresponding 
diminution  of  product. 

3.  How  many  mols  of  bleaching  powder  (assuming  it 
to  be  the  pure  compound  whose  formula  is  given  in  Ques- 
tion i)  would  be   necessary  to  convert  i  mol  of  potassium 
chloride  into  chlorate?     Calculate  the  weight  of  this  sub- 
stance that  would   react  with  the   20  grams  of   potassium 


POTASSIUM    BROMATE    AND    BROMIDE  153 

chloride  taken,  and  compare  this  amount  with  the  amount 
of  bleaching  powder  actually  taken.  How  much  hydrochlo- 
ric acid  would  it  be  necessary  to  use  if  the  bleaching  powder 
were  actually  this  pure  substance  ? 

4.  The  modern  commercial  method  of  making  potas- 
sium chlorate  is  by  the  electrolysis  of  a  potassium  chloride 
solution.  What  are  the  primary  products  formed  at  the  two 
electrodes  ?  Explain  how  the  secondary  reactions  are  sim- 
ilar to  those  outlined  in  the  introductory  discussion. 


46.     POTASSIUM    BROMATE    AND    POTASSIUM 
BROMIDE 

The  reaction  of  bromine  on  solutions  of  caustic  alkalies 
is  almost  identical  to  that  of  chlorine,  and  in  this  connection 
the  discussion  of  the  preparation  Potassium  Chlorate,  No.  45, 
should  be  read.  The  reaction  of  bromine  on  ammonium 
hydroxide  should  also  be  referred  to  under  the  discussion 
of  Ammonium  Bromide,  No.  5. 

Bromine  itself  is  used  in  this  preparation,  as  it  is  not  so 
difficult  or  disagreeable  to  handle  as  chlorine.  By  its  action 
on  concentrated  potassium  hydroxide  solution  a  great  deal 
of  heat  is  produced,  and  in  this  hot  solution  any  hypobromite 
at  first  formed  is  rapidly  converted  into  bromate,  so  that  as 
a  final  result  i  molecule  of  potassium  bromate  to  5  molecules 
of  potassium  bromide  is  obtained.  By  taking  advantage  of 
the  great  difference  in  the  solubility  of  these  salts,  the  former 
may  be  crystallized  pure  from  the  solution  while  the  mother 
liquor  contains  all  of  the  latter,  in  addition  to  the  small 
amount  of  bromate  that  is  soluble.  The  potassium  bromide 
could  not  well  be  crystallized  pure  from  this  solution,  but 
it  is  possible  to  reduce  the  bromate  present  to  bromide  by 
heating  with  charcoal  and  then  to  crystallize  pure  potassium 
bromide. 


154  NON-METALLIC    ELEMENTS 

Procedure.  —  Dissolve  3 1  grams  of  potassium  hydroxide 
in  31  grams  of  water  in  a  250  cc.  Erlenmeyer  flask.  Place 
40  grams  of  bromine  (12^  cc.)  in  a  small  separatory  funnel, 
and  clamp  the  latter  firmly  in  a  vertical  position.  Place  the 
flask  in  a  pan  of  cold  water,  and  lower  the  stem  of  the  sepa- 
ratory funnel  into  the  flask  until  it  nearly  reaches  the  surface 
of  the  solution.  The  funnel  should  now  be  fastened  rigidly, 
and  the  flask  should  be  floating  on  the  surface  of  the  bath, 
so  that  it  may  be  held  by  the  hand  and  constantly  rotated. 
Open  the  stopcock  of  the  funnel  cautiously,  and  allow  the 
bromine  to  run  into  the  solution  at  the  rate  of  2  or  3  small 
drops  per  second.  The  solution  should  grow  hot,  but  if  the 
reaction  becomes  violent  and  red  vapors  escape  from  the 
flask,  stop  the  flow  of  bromine  for  a  few  moments.  The  reac- 
tion is  complete  when  the  solution  has  acquired  a  permanent 
reddish  yellow  tint,  due  to  a  small  excess  of  bromine.  Cool 
the  solution  completely,  collect  the  crystals  of  potassium 
bromate  on  a  filter,  and  recrystallize  them  once  or  twice 
from  a  small  amount  of  hot  water  until  free  from  bromide 
(test  with  silver  nitrate).  Combine  all  of  the  mother  liquors, 
evaporate  until  a  pasty  mass  is  obtained,  mix  this  thoroughly 
with  5  grams  of  powdered  charcoal,  and  dry  the  mass  com- 
pletely. Pulverize  the  dry  mixture  in  a  mortar,  and  then 
heat  it  to  redness  for  an  hour  in  a  large  porcelain  crucible 
surrounded  with  a  funnel  of  asbestos.  Extract  the  product 
with  60  cc.  of  hot  water,  filter,  wash  the  residue  and  the 
filter  with  an  additional  15  cc.  of  hot  water,  and  evaporate 
the  solution  to  obtain  crystals  of  potassium  bromide. 

Questions 

1.  Write    the    reaction    between    bromine    and    sodium 
hydroxide  in  a  cold  dilute  solution;  in  a  hot  concentrated 
solution. 

2.  Why  is  it  impossible  to  obtain  a  mixture  of  5   mole- 


POTASSIUM    IODATE  155 

cules  of  ammonium  bromide  and  i  molecule  of  ammonium 
bromate  by  the  action  of  bromine  on  a  warm  solution  of 
ammonium  hydroxide  ? 

3.  Test  the  Purity  of  the  Potassium  Bromide.  —  Dissolve 
some  of  the  salt  in  water,  and  acidify  the  solution  with  sul- 
phuric acid.     Appearance  of  the  red  color  of  free  bromine 
indicates  that  the  bromate  was  not  all  decomposed  by  the 
heating  with  charcoal.     Explain  and  give  reactions. 

4.  State  reasons  why  it  would  not  be  feasible  to  purify 
the  by-product,  potassium  bromide,  by  recrystallization  with- 
out first  decomposing  the  bromate. 

47.     POTASSIUM    IODATE 

As  is  well  known,  the  chemical  affinity  of  the  halogens 
for  hydrogen  or  positive  elements  decreases  in  passing  from 
fluorine  to  iodine ;  but  the  affinity  for  oxygen  increases  in 
this  order,  so  that  iodates  and  iodic  acid  (I2O5)  are  much 
more  stable  than  chlorates  and  chloric  acid  (C12O5).  Use 
is  made  of  this  fact  in  the  following  preparation,  in  which 
the  total  change  is  represented  fairly  closely  by  the  equation, 

KC1O3  +  I  ->  KIO3  +  Cl. 

The  actual  reaction,  however,  is  not  so  simple  as  this.  The 
presence  of  a  small  amount  of  acid  is  necessary  to  make  it 
take  place.  This  acid  gives  rise  to  a  little  free  chloric  acid, 
which  is  a  far  stronger  oxidizing  agent  than  potassium  chlo- 
rate, and  oxidizes  the  iodine  to  iodic  acid  (I2O5).  By  this 
reaction  more  acid  (HIO8  or  HC1)  is  generated,  and  thus 
the  reaction  when  once  started  proceeds  to  completion.  It 
will  be  noticed  that  in  carrying  out  the  following  directions 
more  iodine  is  taken  than  is  necessary  to  react  with  the 
potassium  chlorate  according  to  the  equation  given  above. 
This  excess  of  iodine  is  oxidized  to  iodic  acid  by  a  part  of 


156  NON -METALLIC  ELEMENTS 

the  free  chlorine  which  is  represented  in   the  equation  as 
escaping. 

Procedure.  —  Dissolve  30  grams  of  potassium  chlorate 
by  warming  it  with  100  cc.  of  water  in  an  800  cc.  flask. 
Add  35  grams  of  powdered  iodine  and  hang  a  small  funnel 
in  the  neck  of  the  flask  to  prevent,  to  some  extent,  the 
escape  of  iodine  vapor.  Place  a  pan  of  cold  water  close 
at  hand ;  then  add  i  cc.  of  nitric  acid  (1.2)  to  the  flask,  and 
warm  rather  carefully  until  a  brisk  reaction  commences. 
Then  allow  the  reaction  to  proceed  so  that  violet  vapors 
fill  the  flask,  but  no  appreciable  quantity  of  iodine  escapes 
through  the  funnel.  If  the  reaction  grows  more  violent  than 
this,  check  it  by  dipping  the  flask  for  a  moment  in  the  cold 
water.  When  the  reaction  is  complete,  boil  the  solution 
until  the  last  trace  of  iodine  has  disappeared.  Then  add 
i  gram  more  of  iodine,  and  boil,  first  in  the  flask  and  then 
in  a  beaker,  until  the  odor  of  chlorine  can  no  longer  be 
detected.  The  solution  now  contains  a  considerable  quan- 
tity of  iodic  acid  in  addition  to  the  potassium  iodate.  Add 
a  solution  of  potassium  hydroxide  until  the  neutral  point  is 
just  reached  (test  by  dipping  a  stirring  rod  in  the  solution 
and  touching  it  to  litmus  paper).  Allow  the  solutior  to 
cool,  collect  the  crystals  of  potassium  iodate,  and  evaporate 
the  mother  liquor  to  obtain  another  crop  of  crystals.  Purify 
the  entire  product  by  recrystallizing  once  from  hot  water. 

Questions 

i .  Experiment.  —  To  3  drops  of  potassium  iodide  solu- 
tion in  10  cc.  of  water  add  freshly  prepared  chlorine  water, 
drop  by  drop,  until  the  iodine  color  which  at  first  appears 
has  been  bleached.  What  change  takes  place  in  the  state 
of  oxidation  of  the  iodine,  first  when  it  is  liberated  from  the 
potassium  iodide,  and  second  when  it  is  further  oxidized  to 
iodic  acid  ?  Write  the  equation  for  the  latter  action, 


IODIC    ACID  157 

2.  Dissolve  a  few  small  crystals  (0.05  gram)  of  potas- 
sium iodate  in  3  cc.  of  warm  water,  and  add  sulphurous 
acid,  drop  by  drop,  to  this  solution,  noting  the  successive 
changes  that  occur  until  the  solution  again  becomes  clear 
and  colorless.  Trace  the  changes  in  the  state  of  oxidation 
of  the  iodine,  giving  reactions,  and  compare  with  the  changes 
in  Experiment  i. 


48.     IODIC   ACID;  IODINE    PENTOXIDE 

Iodine  pentoxide  is  a  white  solid  substance  that  at 
ordinary  temperatures  is  entirely  stable.  It  cannot  be  pre- 
pared by  direct  synthesis  from  iodine  and  oxygen,  because 
when  cold  the  elements  combine  too  slowly,  and  when 
heated  the  compound  is  decomposed  into  the  elements.  It 
may  be  readily  prepared  by  the  direct  oxidation  of  iodine 
by  means  of  strong  oxidizing  agents,  such  as  concentrated 
nitric  acid  or  chlorine.  One  method  for  the  oxidation  of 
iodine  has  already  been  illustrated  under  the  preparation 
of  Potassium  Iodate,  No.  47,  but  there  the  conditions  were 
such  that  a  salt  of  iodic  aoid  was  obtained  rather  than  the 
free  acid  or  its  anhydride.  Starting  with  this  salt,  however, 
the  free  acid  is  easily  obtained  by  metathetical  reactions 
which  depend  on  the  insolubility  of  barium  iodate  and  the 
still  greater  insolubility  of  barium  sulphate. 

Procedure.  —  Dissolve  43  grams  of  potassium  iodate  and 
26  grams  of  barium  nitrate,  separately,  each  in  250  cc.  of 
hot  water,  and  mix  the  two  solutions  at  the  boiling  tempera- 
ture while  stirring  well.  Cool  the  mixture,  let  the  heavy 
precipitate  settle,  decant  off  the  clear  liquid,  and  wash  the 
salt  twice  by  decantation  with  pure  water.  Drain  the 
barium  iodate  on  a  Witt  filter,  and  wash  it  on  the  filter 
with  cold  water.  Then  remove  it  to  a  porcelain  dish,  sus- 
pend it  in  250  cc.  of  water,  heat  to  boiling,  and  stir  in  a 


158  NON-METALLIC    ELEMENTS 

solution  of  15  grams  of  concentrated  sulphuric  acid  (8  cc.) 
in  100  cc.  of  water.  Keep  this  mixture  well  stirred  at  the 
boiling  temperature  for  at  least  10  minutes,  since  the  con- 
version of  solid  barium  iodate  into  solid  barium  sulphate 
is  a  reaction  that  requires  some  time.  Filter  the  solution 
and  rinse  the  last  of  the  iodic  acid  from  the  solid  barium 
sulphate  by  washing  two  or  three  times  on  the  filter  with 
small  portions  of  water.  Evaporate  the  solution  in  a  casse- 
role to  a  small  volume,  and  finally,  holding  the  casserole  in 
the  hand,  keep  the  contents  rotating,  so  that  the  whole 
inside  of  the  dish  is  continually  wet,  and  evaporate  until 
solid  iodic  acid  separates  in  some  quantity.  Cool  com- 
pletely and  rinse  the  crystals  with  three  successive  portions 
of  10  cc.  each  of  concentrated  nitric  acid  (sp.  gr.  1.42),  tritu- 
rating the  crystals  thoroughly  with  each  portion  of  the  acid. 
Warm  the  casserole  carefully  until  the  product  is  perfectly 
dry  and  ceases  to  give  off  acid  vapors.  This  warming  will 
convert  the  iodic  acid  to  a  large  extent  into  the  anhydride 
I2O5.  Place  the  iodine  pentoxide  at  once  in  a  sample 
bottle  or  tube.  . 

To  obtain  surely  anhydrous  iodine  pentoxide,  the  prod- 
uct could  be  heated  for  some  time  in  an  oven  at  about  200°. 
Crystallized  iodic  acid  could  be  obtained  by  dissolving  the 
product  in  a  very  little  water,  in  which  it  is  extremely 
soluble,  and  allowing  the  solution  to  evaporate  slowly. 

Questions 

1.  Experiment.  —  Dissolve  a   little  of   the  iodine  pen- 
toxide in  water.     Test  the  solution  in  such  a  manner  as  to 
show  whether  it  contains  a  strong  acid.     NOTE  :  A  test  with 
litmus  is  not  conclusive,  for  the  preparation  may  still  con- 
tain a  trace  of  nitric  or  sulphuric  acid  which  has  not  been 
completely  removed. 

2.  Experiment.  —  Heat  \  gram  of  iodine  pentoxide  in 


POTASSIUM    PERCHLORATE  159 

a  dry  test  tube.  Insert  a  glowing  splinter  in  the  tube. 
Note  whether  the  entire  substance  can  be  volatilized ;  also 
if  any  of  the  original  substance  deposits  in  the  cooler  part 
of  the  tube. 


49.     POTASSIUM    PERCHLORATE 

When  potassium  chlorate  is  heated  to  about  400°  it  may 
decompose  according  to  either  of  the  following  independent 
reactions : 

4KC1O3  =  KC1  +  3KC1O4,  (i) 

KC1O3    =  KC1  -f  3O.  (2) 

The  second  reaction  is  accelerated  by  catalyzers,  such  as 
manganese  dioxide  or  ferric  oxide,  or  in  fact  any  material 
with  a  rough  surface.  Too  high  a  temperature  also  causes 
reaction  (2)  principally  to  take  place.  On  the  other  hand, 
if  the  temperature  is  maintained  at  the  right  point,  and  the 
salt  is  free  from  dirt,  and  the  inside  of  the  crucible  is  per- 
fectly clean  and  free  from  roughness,  the  decomposition 
proceeds  mainly  according  to  reaction  (i).  Potassium  per- 
chlorate,  being  very  sparingly  soluble  in  cold  water,  may 
easily  be  separated  from  potassium  chloride  and  any  unde- 
composed  potassium  chlorate  by  means  of  crystallization. 

Procedure.  —  Place  50  grams  of  potassium  chlorate  in  a 
dry,  clean  100  cc.  porcelain  crucible,  the  glaze  of  which  is 
in  perfect  condition.  Place  a  small  watch  glass  over  the 
crucible  to  prevent  loss  of  particles  of  the  salt  by  decrep- 
itation, and  heat  gently  until  the  charge  just  melts.  Then 
remove  the  watch  glass  and  keep  the  melt  just  hot  enough 
to  keep  up  a  gentle  evolution  of  oxygen,  but  do  not  increase 
the  temperature  when  the  mass  shows  a  tendency  to  grow 
solid.  At  the  end  of  about  20  minutes  the  melt  should 
begin  to  solidify  around  the  edges  and  should  become  more 


l6o  NON-METALLIC    ELEMENTS 

or  less  pasty  or  semi-solid  throughout;  when  this  point  is 
reached,  let  the  contents  of  the  crucible  cool  completely, 
then  cover  it  with  50  cc.  of  water,  and  let  it  stand  until  it 
is  entirely  disintegrated.  Collect  the  undissolved  potassium 
perchlorate  on  a  Witt  filter  and  wash  it  with  two  successive 
portions  of  15  cc.  of  cold  water  (see  Note  5  (a)  on  page  12). 
Redissolve  the  salt  in  hot  water  (see  solubility  table)  and 
allow  it  to  recrystallize.  About  30  grams  of  potassium  per- 
chlorate should  be  obtained.  A  few  crystals  of  the  product 
should  give  no  yellow  color  (C1O2)  when  treated  with  a 
few  drops  of  concentrated  hydrochloric  acid.  The  product 
should  be  entirely  free  from  chloride  (test  with  silver 

nitrate). 

Questions 

1.  Why  is  manganese  dioxide  added  when  oxygen  is 
prepared  by  heating  potassium  chlorate  ? 

2.  What  is  the  reaction  of  hydrochloric  acid  with  hypo- 
chlorous,  chloric,  and  perchloric  acids,  respectively  ? 

3.  What  are  the  four  oxyacids  of  chlorine?     Compare 
their  stability. 

4.  To  what  extent  are  hydrochloric,  hypochlorous,  chlo- 
ric, and  perchloric  acids  dissociated  electrolytically  in  dilute 
solution  ? 

5.  How  could  pure  perchloric  acid  be  prepared  from 
potassium  perchlorate? 

6.  What  is  the  solubility  of  silver  chlorate  and  of  silver 
perchlorate  ?     How  may  preparations  of  chlorates  and  per- 
chlorates  be  tested  for  the  presence  of  chlorides  ? 

50.     SODIUM    THIOSULPHATE   (Na2S2O8.5H2O) 

Sodium  sulphite  is  a  salt  of  the  lower  oxide  of  sulphur, 
and  may  thus  be  regarded  as  unsaturated  with  respect  to 
oxygen  ;  it  is,  in  fact,  capable  of  slowly  absorbing  oxygen 


SODIUM    THIOSULPHATE  l6l 

from  the  air  and  thereby  going  over  into  sulphate.  If 
it  is  allowed  to  react  with  sulphur,  the  latter  enters  into 
the  compound  in  much  the  same  way  as  does  oxygen,  and 
//&z'0sulphate  instead  of  sulphate  is  formed.  The  sulphur 
so  taken  up  certainly  plays  a  different  function  from  the 
sulphur  already  contained  in  the  compound,  although  it 
is  perhaps  a  question  whether  the  thiosulphate  is  exactly 
the  same  compound  as  sulphate,  except  that  one  oxygen 
atom  is  replaced  by  a  sulphur. 

Sodium  sulphite  is  conveniently  prepared  by  allowing 
sulphur  dioxide  (sulphurous  acid)  to  react  with  sodium 
carbonate.  It  is  practically  impossible,  however,  to  dis- 
tinguish the  exact  point  at  which  the  normal  sulphite 
(Na2SO3)  is  formed;  therefore  it  is  more  expedient  to  divide 
a  given  amount  of  sodium  carbonate  into  two  equal  parts, 
to  fully  saturate  one  part  with  sulphur  dioxide,  whereby 
sodium  bisulphite,  NaHSO8,  is  formed,  and  to  add  the  other 
half  of  the  sodium  carbonate,  thereby  obtaining  the  normal 
sulphite,  Na2SO3. 

Procedure.  —  Dissolve  100  grams  of  sodium  carbonate 
(anhydrous)  in  300  cc.  of  hot  water,  and  divide  the  solution 
into  two  equal  parts.  Reserve  one  part  and  place  about 
five-sixths  of  the  other  half  in  one  flask  and  the  remainder 
in  another  flask.  Connect  these  flasks  in  series  so  that 
sulphur  dioxide  gas  may  be  passed  first  into  the  larger 
volume  of  solution,  and  what  is  there  unabsorbed  may  pass 
on  through  the  second  flask.  Draw  the  gas  from  a  steel 
cylinder  of  liquefied  sulphur  dioxide,  if  one  is  available, 
otherwise  generate  it  by  the  action  of  copper  with  concen- 
trated sulphuric  acid,  and  pass  a  vigorous  stream  of  the 
gas  into  the  solutions.  After  a  short  time  a  marked  frothing 
occurs  in  the  first  flask,  due  to  the  escape  of  carbon  dioxide, 
and  after  this  frothing  ceases  a  similar  frothing  soon  com- 
mences in  the  second  flask.  When  the  latter  ceases,  pass 


1 62  NON-METALLIC    ELEMENTS 

the  gas  a  little  while  longer  until  sulphur  dioxide  escapes 
freely  from  the  second  bottle.  Then  place  the  solution  of 
sodium  bisulphite  in  a  600  cc.  beaker,  and  stir  in  rather 
slowly  the  remaining  sodium  carbonate.  Add  45  grams  of 
flowers  of  sulphur,  cover  the  beaker  with  a  watch  glass, 
and  keep  the  mixture  just  barely  boiling  for  an  hour  or 
longer.  Filter  the  solution,  concentrate  it  to  a  volume  of 
about  200  cc.,  and  leave  it  uncovered  over  night  to  crystal- 
lize in  a  place  free  from  dust.  Collect  the  crystals  and 
obtain  further  crystals  from  the  mother  liquor. 

Questions 

1.  Experiment.  —  Dissolve  \  gram  of   the  product  in 
5  cc.  of  water  and  add  2  cc.  of  hydrochloric  acid.     Observe 
the  odor  and  the  precipitate.     What  is  the  free  acid  corre- 
sponding to  the  salt,  sodium  thiosulphate  ?     What  can  be 
said  regarding  the  stability  of  this  acid? 

2.  What  is  the  valence  of  sulphur  in  each  of  the  salts, 
sodium   sulphide,  sodium    sulphite,    and   sodium   sulphate? 
State  in  each  case  whether  the  sulphur  plays  the  part  of 
a  positive  or  negative  element. 

3.  Distinguish  between   the   parts  played   by  the  two 
atoms  of  sulphur  in  sodium  thiosulphate. 

4.  Give  equations  to  represent  the  successive  reactions 
that  take  place  when  sulphur  dioxide  is  passed  into  a  sodium 
carbonate  solution.     What  stage  of  the  process  is  indicated 
by  each  of  the  succeeding  phenomena  ?     (a)  The  gas  passes 
into  the  solution  in   distinct  bubbles  and  is  in  large  part 
absorbed,     (b)  Effervescence  takes  place  with  minute  bub- 
bles arising  from  every  part  of  the  solution,     (c)  Efferves- 
cence ceases,  and  the-  gas  enters  the  solution  again  in  clear, 
distinct  bubbles,  but  still  it  is  for  the  most  part  absorbed. 
(d]  The  gas  passes  through  the  solution  in  distinct  bubbles 
and  is  entirely  unabsorbed. 


GENERAL    QUESTIONS    VII  163 

GENERAL    QUESTIONS.     VII 

NON-METALLIC    ELEMENTS   OF   THE    SIXTH    AND    SEVENTH    GROUPS 
OF   THE   PERIODIC   SYSTEM 

i.  What  is  the  valence  of  oxygen  and  sulphur  towards 
hydrogen  and  metallic  elements?  What  is  the  valence  of 
the  halogens? 

?.  Give  the  symbols  of  the  oxides  of  sulphur  and  of 
the  halogens  which  can  actually  be  made.  State  which  of 
them  are  salt-forming  oxides.  Give  the  symbols  of  the  most 
important  oxyacids  of  sulphur,  chlorine,  bromine,  and  iodine, 
and  state  in  each  case  what  oxide  (actual  or  hypothetical) 
is  to  be  regarded  as  its  anhydride. 

3.  Give  data  regarding  the  ease  with  which  the  hydro- 
gen compounds  of  oxygen,  sulphur,  and  the  halogens  can 
be    formed ;    also   state    how   readily  these   compounds    are 
decomposed  by  heat.     Draw  conclusions  as  to  the  relative 
chemical  activity  of  these  elements  when  they  act  as  negative 
elements. 

4.  Compare  these  hydrogen  compounds  with  regard  to 
their  acid  strength  when  they  are  dissolved  in  water. 

5.  Compare  the  same  elements  with  regard  to  the  ease 
with   which   they   form   compounds    with    oxygen,   and   the 
ease   with   which    these   compounds   can    be    decomposed. 
Draw  conclusions  as  regards  the  chemical  activity  of  these 
elements  when  they  act  as  positive  elements. 

6.  Compare  the  degree  of  electrolytic  dissociation   in 
aqueous  solution  of  the  acids  derived  from  the  oxides,  SO2, 
SO3,   C12O,   C12O5,   C12O7.     When    an   element   forms    more 
than   one   oxide,   which,   as   a   general   rule,   has   the   more 
pronounced  acid  character? 


APPENDIX 

ADDITIONAL   GENERAL    QUESTIONS 

I.     ALKALI  AND  ALKALINE  EARTH  METALS 

1.  What  are  the  symbols  of  the  oxides  of  sodium  and 
potassium  in  which  these  metals  undoubtedly  display  their 
ordinary  valence  of  one  ?     How  can   these  oxides  be  pre- 
pared, using  the  peroxides  as  a  starting  point  ?     How  can 
they  be  prepared  from  the  hydroxides  or  from  the  nitrates  ? 

2.  What  oxide  is  obtained  by  burning  sodium  in  the 
air?     What  relation  does  this  oxide  bear  to  hydrogen  perox- 
ide?    What  is  supposed  to   be   the  molecular  structure  of 
a  peroxide  ? 

3.  What   is   the  difference  in  meaning   between   the 
terms  base  and  alkali? 

4.  Distinguish   between    a   mild  alkali    and   a   caustic 
alkali. 

5.  Explain  why  solutions  of  the  carbonates  of  sodium 
and  potassium  behave  as  mild  alkalies. 

6.  Judged  by  its  composition,  sodium  bicarbonate  is 
an  acid  salt.     Explain  why  in  spite  of  this  fact  its  solution 
behaves  as  a  mild  alkali. 

7.  What  is  the  difference  between  ammonia  and  am- 
monium ?     How  can  ammonia  be  set  free  from  ammonium 
salts  ? 

8.  In    its    compounds   the    ammonium    radical,    NH4, 
behaves  similarly  to  the  metals  of  the  alkali  group.     It  is 
not  unreasonable  to  suppose  that  if  ammonium  could  be 
isolated,  it  would  show  properties  similar  to  those  of  sodium 

165 


1 66  APPENDIX 

and  potassium  when  in  the  metallic  state.     To  what  extent 
is  this  supposition  borne  out  by  the  facts  ? 

9.  Why  is  the  alkaline  strength  of  an  ammonium 
hydroxide  solution  greatly  reduced  by  the  addition  of  an 
ammonium  salt  ? 

10.  Give  an  outline  of  the  modern  method  for  preparing 
the  alkali  metals  and  the  alkaline  earth  metals.     What  are 
the  most  characteristic  properties  of  these  metals  ? 

11.  What  are  the  chief  differences  between  the  prop- 
erties  of   the    alkali    and    alkaline    earth    metals    (compare 
potassium  with  calcium)  ? 

12.  Tabulate   the    solubilities   of   the   hydroxides,   car- 
bonates, and    sulphates  of  the   alkaline   earth  metals,   and 
observe  in  what  sense  the  solubility  changes  with  increasing 
atomic  weight  of  the  metal  in  each  series  of  salts. 

II.     ELEMENTS  OF  THE  THIRD  GROUP  OF  THE 
PERIODIC  SYSTEM 

1.  Give  the   formulae  of    normal   boric   and   aluminic 
acids;  of  the  meta  acids;  of  the  anhydrides;  and  of  the 
sodium    salts    of   tetraboric    acid,   aluminic    acid,   and   met- 
aluminic  acid. 

2.  All  of  the  alums  are  isomorphous.     What  is  iso- 
morphism ?     Give  the  name  and   symbol   of   at   least  two 
alums  which  contain  neither  aluminum  nor  potassium. 

It  is  not  possible  to  prepare  an  alum  from  boron  in 
which  this  element  plays  the  part  of  aluminum  in  common 
alum.  What  is  the  fundamental  difference  between  boron 
and  aluminum  which  accounts  for  this  difference  in  behavior  ? 

3.  Before   the   use   of   the   electrolytic  process,  what 
metal  was  used  to  reduce  the  oxide  of  aluminum  in  the 
preparation  of  metallic  aluminum? 

4.  Describe    the    principle    of    the    alumino-thermic 
processes. 


APPENDIX  167 

5.  Why  would  it  not  be  possible  to  prepare  metallic 
calcium  by  means  of  the  aluinino-thermic  process  ? 

6.  What  other  metals  could  be  substituted  for  aluminum 
in  the  alumino-thermic  process  ? 

7.  Experiment. — To  a  solution  of  any  aluminum  salt 
(use  common  alum),  add  a  solution  of  sodium  carbonate. 

What  is  the  gas  evolved,  and  what  is  the  precipitate  ? 
Why  is  it  impossible  for  aluminum  carbonate  to  form  under 
these  conditions  ? 


III.     HEAVY  METALS  OF  THE  FIRST  Two  GROUPS 
OF  THE  PERIODIC  SYSTEM 

1.  Cite  facts  which   show  the  difference   in  chemical 
activity  between  the  elements  of  Family  A  and  of  Family  B 
in  Group  I. 

2.  Compare  the  chlorides  of  copper,  silver,  and  gold 
(copper  and  gold   each  possess  two  chlorides)  with  regard 
to  their  solubility  and  the  valence  of  the  metal.     Which  of 
the   chlorides   are   characteristic   of   the   position   of   these 
elements  in  Group  I  ? 

3.  Starting  with  an  alloy  of  copper  and  silver,  devise 
a  method  for  obtaining  the  two   metals  separately  in   the 
metallic  condition. 

4.  Experiment.  —  Prepare  cuprous  oxide  by  reducing 
a  cupric  salt  with  grape  sugar  in  an  alkaline  solution.     To 
20  cc.  of  a  10  per  cent,  solution  of  copper  sulphate  add 
50  cc.  of  water,  and  then  stir  in  20  cc.  of  a  10  per  cent, 
solution  of  sodium  hydroxide ;  add  a  solution  of  grape  sugar 
until  a  deep  blue  color  appears  and  the  precipitate  of  cupric 
hydroxide  has  nearly  or  quite  dissolved;   lastly,  warm  the 
mixture  gently  until  a  yellow  and  then  a  clear,  bright  red 
precipitate  is  obtained.     Collect  this  precipitate  of  cuprous 
oxide  on  a  filter  and  wash  it  with  water.     Treat  the  cuprous 


I 68  APPENDIX 

oxide  on  the  filter  with  dilute  sulphuric  acid,  and  examine 
the  solution  formed  and  the  insoluble  residue  to  determine 
what  products  are  formed  in  the  reaction.  Explain  how 
one-half  of  the  cuprous  oxide  is  oxidized  at  the  expense  of 
the  other  half  which  is  reduced. 

5.  Experiment.  —  Add  a  little  cupric  salt  solution  to 
a  potassium  iodide  solution,  and  test  a  drop  of  the  resulting 
liquid  for  the  presence  of  free  iodine.     What  is  the  precipi- 
tate ?     Rinse  it  free  from  the  dark-colored  solution  in  order 
to  observe  its  own  color.     Write  the  reaction,  and  explain 
how  the  copper  salt  has  acted  as  an  oxidizing  agent. 

6.  Compare  the  chlorides  of  zinc,  cadmium,  and  mer- 
cury (mercury  has  two  chlorides)  as  regards  their  solubility 
and  the  valence  of  the  metal. 

7.  In  which  of  their  compounds  do  zinc,  cadmium, 
and  mercury  most  resemble  each  other,   and   thus  justify 
their  classification  in  the  same  family  of  the  periodic  system  ? 

8.  In  what  important  respects  do  these  metals  differ 
from  the  alkaline  earth  metals  which  form  the  other  family 
in  the  second  group  of  the  elements  ? 

9.  Given  an  alloy  of  zinc  cadmium  and  mercury,  how 
might  one  proceed  »to  obtain  the   three  separately  in  the 
metallic  condition  ? 

i  o.  Experiment.  —  Prepare  a  dilute  acidified  solution  of 
zinc  sulphate,  using  10  cc.  of  a  10  per  cent,  solution  and 
adding  20  cc.  of  water  and  5  cc.  of  sulphuric  acid  (i  :  4). 
Divide  the  solution  into  two  parts,  and  treat  one  part  with 
hydrogen  sulphide  and  the  other  part  with  ammonium 
sulphide. 

Explain  why  the  solubility  product  of  zinc  sulphide  is 
exceeded  in  the  one  case  and  not  in  the  other. 

n.  Experiment.  —  Saturate  with  hydrogen  sulphide  an 
unacidified  solution  of  zinc  sulphate  (made  by  diluting  5  cc. 
of  the  10  per  cent,  solution  with  10  cc.  of  water),  when 


APPENDIX  169 

saturated  the  solution  will  continue  to  smell  of  hydrogen 
sulphide  after  closing  the  mouth  of  the  tube  with  the  thumb 
and  shaking.  Filter  off  the  precipitate,  and  test  for  zinc 
ions  in  the  filtrate  by  adding  ammonium  sulphide. 

Explain  the  cause  of  the  incompleteness  of  this  reaction. 

12.  Experiment.  —  Modify  the  last  experiment  by  dis- 
solving 3  grams  of   solid  sodium   acetate   in   the   solution 
before    saturating   with    hydrogen    sulphide.     Do    any    zinc 
ions  pass  into  the  filtrate  in  this  case  ?     How  is  the  differ- 
ence in  the  results  explained  ? 

13.  Experiment.  —  Dilute  5  cc.  of  a  10  per  cent,  solution 
of  cadmium  chloride  or  cadmium   sulphate  with    10  cc.  of 
water,  and  saturate  with  hydrogen  sulphide.     Filter  off  the 
precipitate,  and  test  for  cadmium    ions    in   the   filtrate    by 
adding  ammonium  sulphide. 

Saturate  with  hydrogen  sulphide  a  similar  solution  of 
cadmium  salt  after  first  adding  10  cc.  of  hydrochloric  acid 
(1.12  sp.  gr.). 

How  do  the  solubility  products  of  cadmium  and  zinc  sul- 
phides compare  with  each  other  ? 

14.  Find   out   from    the    text-book,  or    by   experiment, 
whether  the  sulphides  of  mercury,  copper,  and  silver  can 
be  precipitated  from  moderately  acidified  solutions.     Com- 
pare these  sulphides  with  those  of  zinc  and  cadmium  with 
respect  to  their  solubility  products. 

15.  Solutions  of  mercuric  chloride   and   mercuric  cya- 
nide are  poor  conductors  of  electricity.     To  what  general 
rule  do  these  salts  thus  form  an  exception  ? 

IV.     ELEMENTS  OF  THE  FOURTH  GROUP  OF  THE 
PERIODIC  SYSTEM 

i.  Why  is  carbon  of  importance  in  the  organic  world  ? 
With  what  other  elements  is  carbon  found  in  combination 
in  organic  compounds  ? 


I/O  APPENDIX 

2.  In  what  compounds  does  carbon   most  frequently 
occur  in  rocks  ? 

3.  What  is  the  relation  between  an  ortho  and  a  meta 
acid  ?     Give  symbols  of  ortho  and  meta  silicic,  stannic,  and 
plumbic  acids. 

4.  Experiment.  —  Add  hydrochloric  acid  to  a  solution 
of   sodium  silicate   (water  glass).     What   is  the  gelatinous 
substance  obtained,  and   what  change  would   it  undergo  if 
dried  and  baked  ? 

5.  Name  some  of  the  common  minerals  that  consist 
of  some  form  of  silicic  acid  or  its  anhydride. 

6.  With  what   bases   is  silicic   acid   most   commonly 
found  in  combination  in  the  minerals  that  constitute  our 
common  rocks  ? 

7.  Place  tin  and  lead  approximately  in  their  places 
in   the   electromotive   series   of   the    metals.     Explain   why 
either  of  these  metals  will  react  slowly  with  hydrochloric 
acid  but  rapidly  with  nitric  acid. 

8.  By  what  oxidizing  agents  can  a  stannous  compound 
be  converted  into  a  stannic  ?     By  what  means  can  a  stannic 
compound  be  reduced  to  a  stannous  ? 

9.  Under  what  conditions  can   lead  in  the  plumbous 
state  be  oxidized  to  the  plumbic  state?     Is  this  as  easily 
accomplished  as  the  corresponding  oxidation  of  stannous  to 
stannic  tin  ?     Under  what  conditions  does  lead  dioxide  act 
as   an   oxidizing   agent?     (See    Experiment  4  under   Lead 
Dioxide,  page  95.) 

10.  Show  how  the  action  of  the  common  lead  storage 
battery  depends  upon  changes  in  the  state  of  oxidation  of 
lead. 

11.  Show   how  the  mixed  oxides  of  lead,   Pb2O3   and 
Pb3O4,  may  be  regarded  as  salts  of  ortho  and  meta  plumbic 
acid. 

1 2 .  Find  out  the  following  facts  either  from  a  text-book 


APPENDIX  I/I 

or  by  means  of  experiment :  What  is  the  color  and  what  is 
the  chemical  formula  of  the  precipitates  formed  when  hydro- 
gen sulphide  is  passed  into  solutions  of  PbQ2,  SnCl2,  and 
SnCl4,  respectively?  Are  all  of  the  lead  and  tin  thrown  out 
in  this  way  from  faintly  acidified  solutions  ?  Does  precipi- 
tation occur  in  solutions  strongly  acidified  with  hydrochloric 
acid  ? 

Compare  lead  sulphide  and  stannous  sulphide  with  the 
sulphides  of  other  heavy  metals  in  regard  to  their  solubility 
product. 

13.  If  a  precipitate  was  obtained  containing  sulphides 
of  both  tin  and  lead,  how  might  it  be  treated  in  order  to 
separate  the  two  metals  ?  (See  Experiment  2  under  Stannic 
Sulphide,  page  89.) 

V.     ELEMENTS  OF  THE  FIFTH  GROUP  OF  THE 
PERIODIC  SYSTEM 

1.  Describe  the  physical  properties  of  the  pure  trichlo- 
rides of  phosphorus,  arsenic,  and  antimony.     Compare  the 
properties  of  these  compounds  with  those  of  the  tetrachlo- 
rides  of   carbon,   silicon,   and  tin.     Are   such  chlorides    of 
the  metalloids  to  be  regarded  as  salts? 

2.  What  is  the  behavior  of  the  trichlorides  of  phos- 
phorus, arsenic,  antimony,  and  bismuth  when  treated  with 
water  ?     Show  how  their  varying  behavior  corresponds  with 
the  increasing  metallic  properties  of  the  elements. 

3.  Give  reasons  why  arsenic  and  antimony  should  be 
classed  as  metals ;  as  non-metals.     Which  one  of  the  metals 
studied  under  Group  IV  also  falls  into  the  transitional  class 
between  metals  and  non-metals? 

4.  Distinguish  between  ortho,  meta,  and  pyro  phos- 
phoric  acids.     To  which   of   these  acids   does   nitric  acid 
correspond  in  structure  ? 


I72  APPENDIX 

Which  of  the  antimonic  acids  yields  a  sparingly  soluble 
sodium  salt  ? 

5.  Distinguish  between   phosphorous  and   phosphoric 
acids. 

Give  equations  to  show  reactions  in  which  nitrous, 
phosphorous,  and  arsenious  acids  act  as  reducing  agents. 
Give  an  example  of  the  action  of  arsenic  acid  as  an  oxidizing 
agent. 

6.  What  is  meant  by  antimonyl  and  bismuthyl  salts  ? 
Give  symbols  and  describe  their  properties.     By  what  other 
names  are  bismuthyl  nitrate  and  antimonyl  chloride  known  ? 

7.  Arsenic  pentasulphide  can  be  slowly  precipitated 
by  leading  hydrogen  sulphide  into  a  solution  of  arsenic  acid 
containing  a  large  amount  of  concentrated  hydrochloric  acid. 
Does  this  fact  give  any  indication  of  the  extent  to  which  the 
ion  As+H    ++  is  capable  of  existing  in  solution  ? 

VI.     HEAVY  METALS  OF  THE  SIXTH,  SEVENTH,  AND 
EIGHTH  GROUPS  OF  THE  PERIODIC  SYSTEM 

1.  What  is   the  valence   of   the   metal    in   the  oxides 
Mn8O4  and  Fe3O4  ?     Explain  in  what  sense  these  oxides  may 
be  regarded  as  salts. 

2 .  Experiment.  —  Oxidation   of   Manganous   Salts   to 
Permanganate  in  Acid  Solution :  To  half  a  test  tube  of  water 
add  3  drops  of  a  manganous  sulphate  solution  and   10  cc. 
of  nitric  acid.     Then  add  \  gram  of  lead  dioxide  and  boil 
the  mixture.     Let  the  solid  settle,  and  observe  the  color  of 
the  clear  solution. 

Explain  why  the  presence  of  hydrochloric  acid  would 
prevent  the  formation  of  the  red  color.  Recall  the  action 
of  hydrochloric  acid  with  potassium  permanganate. 

3.  Experiment.  —  Oxidation  of  a  Chromic  Salt  in  Alka- 
line Solution:  To   a   little  of   a   chromic  salt   solution   add 


APPENDIX  173 

three  times  its  volume  of  10  per  cent,  sodium  hydroxide 
solution.  Warm  the  mixture  and  pass  in  chlorine  until  the 
color  has  become  a  clear  yellow.  Compare  this  action  with 
the  action  of  hydrochloric  acid  with  sodium  chromate,  and 
explain  how  the  different  conditions  make  these  practically 
opposite  reactions  possible. 

4.  Write  the  reactions  for  the  oxidation   of   a   ferrous 
salt  by  (a)  nitric  acid ;  (b)  bromine ;  (c)  potassium  perman- 
ganate.    Write  the  reactions  for  the  reduction   of  a  ferric 
salt  by   (d)    hydrogen   sulphide;   (e)    sulphur  dioxide;   (/) 
stannous  chloride. 

5.  From  what  oxide  of   iron    are  the   ferrates  derived, 
and   is  this  oxide   acidic  or  basic  in  character  ?     Can  the 
oxide  itself  be  prepared  ?     How  is   potassium  ferrate   pre- 
pared ?     With  what  compounds  of  chromium  and  manganese 
is  it  analogous  in  composition  ? 

6.  Show   that   chromium   exists  in    the  same  state  of 
oxidation    in    both    chromates    and   bichromates.      In    what 
respect  do  these  salts  differ  from  each  other  ? 

Show  that  on  the  other  hand  manganates  and  perman- 
ganates are  derived  from  different  oxides  of  manganese. 

7.  What  salt  is  formed  by  treating  i  mol  of  potassium 
sulphate  with  i  mol  of  sulphuric  acid  ?     How  does  this  salt 
resemble  and  how  differ  from  the  salt  obtained  by  treating 
i  mol  of  potassium  chromate  with  i  mol  of  chromic  acid  ? 

VII.     NON-METALLIC  ELEMENTS  OF  THE  SIXTH  AND 
SEVENTH  GROUPS  OF  THE  PERIODIC  SYSTEM 

i.  The  tension  of  the  elements  chlorine,  bromine,  iodine, 
and  sulphur,  to  pass  into  the  form  of  simple  negative  ions  in 
aqueous  solution,  decreases  in  the  order  given.  Describe 
simple  experiments  which  show  the  correctness  of  this 
statement. 


APPENDIX 

2.  Experiment.  —  Reduction  of  any  Oxysalt  of  Sulphur. 
Mix  a  little  sodium  sulphate  (or  any  oxysalt  of  sulphur)  with 
twice  its  amount  of  sodium  carbonate ;  moisten  the  mixturej 
so  that  some  of  it  may  be  made  to  adhere  to  the  charred  end 
of  a  burnt  match.     Heat  the  mixture  on  the  match  end  in 
the  reducing  part  of  a  Bunsen  flame  until  the   salt  melts. 
Detach  the  end  of  the  match  with  the  fused  salt,  and  place 
it  on  a  bright  silver  coin,  together  with   i    drop  of  water. 
After  a  few  moments   observe  the  dark  fleck  on  the  coin. 
Explain  the  reactions  involved  in  the  test. 

3.  Would  it  be  feasible  to  reduce  sodium  sulphate  with 
charcoal  as  an  industrial  method  of  preparing  sodium  sul- 
phide ?     Discuss  the  relative  advantages  of  this  and  other 
possible  methods. 

4.  If  it  were  desired  to  reduce  potassium  chlorate  to 
chloride,  why  would   it  be  unsafe  to  mix  it  with  charcoal 
and  heat  ?    Why  does  not  the  same  danger  exist  with  sodium 
sulphate  ? 

5.  Experiment.  —  Agitate  a  few  grains  (0.05  gram)  of 
iodine  with  2  cc.  of  water.     To  what  extent  does  the  iodine 
dissolve  ?    Add  a  few  crystals  of  potassium  iodide  (o.i  gram), 
and  observe  whether  any  more  of  the  iodine  passes  into  solu- 
tion.    Account  for  the  increased   solubility  of   the  iodine. 
What  are  polyiodides? 

6.  Experiment.  —  Treat  £  gram  of   powdered  sulphur 
with  5  cc.  of  a  colorless  solution  of  sodium  or  ammonium 
sulphide.     What  is  a  polysulphide  ?      To   a  solution   of   a 
colorless  sulphide  add  hydrochloric  acid  until  the  solution 
is    acid   to    litmus.     Treat   in   the    same  way  some  of  the 
polysulphide  just  prepared  and  compare  the  results. 

7.  When  a  solution  of  potassium  iodide  is  exposed  to 
the  action  of  the  oxygen  of  the  air  no  change  takes  place. 
When,  on  the  other  hand,  a  solution  of  hydriodic  acid   is 
exposed  to  the  same  conditions,  iodine  is  slowly  displaced 


APPENDIX  175 

and  the  solution  becomes  brown,  due  to  the  liberated  iodine. 
Show  by  writing  the  ionic  reactions  that  when  iodine  is  dis- 
placed by  oxygen,  either  hydrogen  ions  must  be  used  up 
or  hydroxyl  ions  must  be  produced,  and  that  therefore  the 
reaction  is  favored  by  the  presence  of  the  former  and  re- 
tarded by  the  presence  of  the  latter. 

8.  Write  reactions  in  which  sulphuric  acid  acts  as  an 
oxidizing  agent ;  in  which  sulphurous  acid  acts  as  a  reducing 
agent;  in  which  sulphurous  acid  acts  as  an  oxidizing  agent; 
in  which  hydrogen  sulphide  acts  as  a  reducing  agent. 

9.  Write    reactions    in    which    hydriodic,   hydrobromic, 
and   hydrochloric   acids   act   as   reducing  agents;   in   which 
the  oxyacids  of  the  halogens  (take,  for  example,  hypochlo- 
rous,  chloric,  and  iodic  acids)  act  as  oxidizing  agents.     Ex- 
plain in  what  way  these  reactions  are  of  the  same  character 
as  these  written  under  Question  8. 


1 76 


APPENDIX 


INTERNATIONAL   ATOMIC    WEIGHTS,    1915 


Symbol. 

Atomic 
weight. 

Symbol. 

Atomic 
weight. 

.   .   .  Al 

27  1 

Mo 

96  0 

.   .   .  Sb 

120.2 

Nd 

1443 

A 

39  9 

Ne 

20  0 

...  As 

74.96 

Nickel     

.   .   .  Ni 

58.68 

-      Ba 

137  37 

N 

14  01 

Be 

9  1 

Os 

190  9 

.   .   .  Bi 

208  0 

o 

16  00 

.   .   .  B 

11  0 

.    .    .  Pd 

106  7 

.   .    .  Br 

79  92 

p 

31  0 

.   .    .  Cd 

112.40 

.  Pt 

1950 

.    .    .  Cs 

132.81 

.   .   .  K 

39  10 

.    .    .Ca 

40.09 

.    .    .  Pr 

1406 

c 

12  00 

Ra 

226  4 

.   .  Ce 

14025 

Rh 

102  9 

Chi     ' 

Cl 

3546 

Rb 

85  45 

mi  m 

Cr 

52.0 

Ru 

101  7 

Cobalt 

Co 

58.97 

Sa 

150  4 

Cb 

93.5 

Sc 

44  1 

Cu 

63.57 

.  Se 

79  2 

Dysprosium   .... 

.    .    .Dy 
Er 

162.5 
167.4 

Silicon     
Silver      

.   .   .Si 
.  Ag 

28.3 

10788 

Eu 

152.0 

.   .   .  Na 

23  00 

Fluorine 

,    .    .  F 

19.0 

.    .    .  Sr 

87.62 

Gd 

1573 

s 

32  07 

.   .   .  Ca 

69.9 

Tantalum    

.    .    .  Ta 

181.0 

Ge 

725 

Tellurium 

Te 

127  5 

Gold     

.    .      Au 

197.2 

Terbium      ..... 

.   .   .Tb 

159.2 

He 

4  0 

Thallium 

Tl 

2040 

H 

1  008 

Th 

232.42 

In 

114.8 

.   .   .Tm 

1685 

I 

126.92 

Tin   

.    .    .  Sn 

119.0 

Ir 

193.1 

.    .    .  Ti 

48.1 

.   .   .  Fe 

55.85 

.    .    .  W 

184.0 

Kr 

83  0 

u 

238  5 

L  y  h       m 

La 

1390 

v 

51  2 

Lead               .   •   • 

.   .  Pb 

207  10 

Xe 

1307 

Lithium  

...  Li 

7.00 

Ytterbium 

Lu 

174.0 

.  Yb 

1720 

Me 

2432 

Yttrium   

.   .   .  Yt 

89.0 

Mn 

54.93 

.   .   .  Zn 

6537 

He 

200.0 

.   .   .  Zr 

90.6 

APPENDIX 


177 


at 
d 


s 


o 


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St 


-Oo? 
02  d 


jap 


21 


«§. 


a 


>-l  I- 
OJ92, 


3 


It 


m 

wl 


a 


TABLE   OF   SOLUBILITIES  i 

In  the  following  tables  are  given  data  which  should  be  useful  in 
connection  with  the  preparations  and  questions  in  this  book. 

The  formulae  given  are  those  of  the  crystallized  compounds  which 
most  readily  separate  from  aqueous  solution  at  the  laboratory  temper- 
ature, but  it  should  be  remembered  that  many  salts  have  several 
hydrates,  and  it  has  often  been  difficult  to  decide  which  one  to  place 
in  the  table. 

In  the  second  column  the  behavior  of  the  crystallized  salt  when  it 
is  exposed  to  the  air  of  the  laboratory  is  indicated :  s  =  unchanged 
by  exposure  to  atmosphere;  e  —  efflorescent;  d  —  deliquescent; 
d,  e  =  deliquescent  or  efflorescent,  according  as  to  whether  the 
humidity  is  above  or  below  the  average ;  COo  =  absorbs  carbon 
dioxide  and  falls  to  a  white  powder ;  Ox  =  compound  is  oxidized, 
especially  in  presence  of  moisture. 

In  the  third  column  are  given  the  figures  for  the  solubility  at 
o°,  25°,  and  100°,  except  in  the  cases  in  which  other  temperatures  are 
indicated  in  parenthesis.  Fractions  have,  as  a  rule,  been  dropped  in 
giving  the  solubilities. 

1  Much  of  the  data  in  this  table  has  been  obtained  from  Seidell,  Solubilities 
of  Inorganic  and  Organic  Substances. 


178 


APPENDIX 


179 


c 

4J 

SOLUBILITY  IN  Wi 

LTER. 

c_u                             Formula  ot 
crystallized  salt. 

Behavior  wh 
exposed  to 
atmosphere 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°                  25°             100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Aluminum  : 
chloride  A1C13  6H2O 

(15°)  70     . 

4 

nitrate,  A1(NO8)3.9H2O    
sulphate,  A12(SO4)3.18H2O  

Ammonium  : 
acetate,  NH4C2H3O2  

..  d  .. 
.  .  s  .  . 

d 

very  soluble 

..31  38  89  .. 

....  0.8 

bromide,  NH4Br  

.  .  s 

(10°)  66     (30°)  81   (100°)  128 

6 

chloride,  NH4C1  
nitrate,  NH4NO3   

Antimony  : 
chloride,  SbCl3  

.  .  s  .  . 

..d  .. 

d 

..29  39  77  .. 
.  118  214  871.. 

}  hydrolyzes  with  water  to 

5 
11 

sulphate,  Sbo(SO4)3  

d 

(insoluble  basic  salt  ; 

sulphide,  SboS3  

.  .  s 

insoluble  ;  soluble  in  concen- 

Arsenic : 
sulphide,  As2S3 

trated  acids 
insoluble  in  water  or  acids  ; 

Barium  : 
carbonate,  BaCO3  
chloride,  BaCl2.2H2O   
chromate,  BaCrO4  . 

.  .  s  .  . 
.  .  s  .  . 

s 

soluble  in  alkalies 

0.0023  
..32  37  59  .. 
00004 

.  .  0.00011 
....  1.7 
0  000015 

hydroxide,  Ba(OH)2.8H2O  
iodate,  Ba(IO3)2.H2O   
nitrate,  Ba(NO3)2  

CO2 

.  .  s  .  . 
.  s,  d  . 

(0°)1.7     (25°)  4.7      (80°)  101 

.0.008....  0.03  0.2.  . 
5              10      .  .   34 

....  0.2 
....0.001 
.  .  03 

sulphate,  BaSO4  
sulohite   BaSO3 

.  .  s  .  . 
s 

0.00023  
(20°)  0  020     (80°)  0  002 

.  .  0.000010 
0  001 

sulphide,  BaS.6H2O  

.  Ox  . 

very  soluble  ;  hydrolyzes  to 

Bismuth  : 
chloride   BiCl3 

d 

Ba(SH)2  and  Ba(OH)2 

nitrate,  Bi(NO3)3.5H2O    
sulphate,  Bi2(SO4)3 

..  d  .. 
d 

(hydrolyzes  with  water  to 
insoluble  basic  salt  ; 
very  soluble  in  acids 

sulphide,  Bi2S3     

.  .  s  .  . 

Cadmium  : 
chloride,  CdCl2.2iH2O  
nitrate   Cd(NO3)2  4H2O 

.  .  e  .  . 
d 

(0°)  90.  (18°)  110.  (ioo°)147 
(0°)110  (18°)  127     (60°)  326 

5 
43 

sulphate,  CdSO4  2f  H2O 

(0°)  76     (40°)  79     (100°)  61 

2 

i8o 


APPENDIX 


c  ,                                   Formula  of 
crystallized  salt. 

Behavior  when 
exposed  to 
atmosphere. 

SOLUBILITY  IN  WATER. 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°               25°              100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Cadmium  : 
sulphide,  CdS  

.  .  S 

insoluble  ;  soluble  in  con- 
centrated acids 

0.0013  
..60  88  159.. 
(18°)  178  

...0.00013 
5.2 
53 

Calcium  : 
carbonate,  CaCO3  
chloride,  CaCl2.6H2O  
chlorate,  Ca(ClO3)2.2H2O    .... 

.  .  S   .  . 

..  d  .. 
..  d 

chromate,  CaCrO4.2H2O  
fluoride  CaF2 

.  .  e  .  . 

g 

..11  12  3... 
00016 

0.8 
0  0002 

hydroxide   Ca(OH)2 

C02 
..  d  .. 

.  0.19....  0.16  ....0.08  . 
(18°)  122 

....  0.02 
52 

nitrate  Ca(NO3)2  4H2O       .... 

oxalate,  CaC2O4.H2O  
sulphate,  CaSO4.2H2O  
sulphite  CaSO3  

.  .  s  .  . 

.  .  s  .  . 

0.0007  
.0.18  ....  0.21  ...  0.16.. 
0  004       

.  .  0.00004 
.  .  .  0.015 
.  .  .0.0003 

Chromium  : 
chloride  CrCH  6H2O 

d 

130 

8 

nitrate  Cr(NO3)s  9H2O 

very  soluble  ;  melts  36.5° 
120  

insoluble  ;  soluble  in  acids 
..42     53  104.  . 

4 
3 

sulphate,  Cr2(SO4)3.18H2O  .... 

Cobalt  : 
carbonate  CoCO3      .... 

.  s,  e  . 
.  .  s  .  . 

chloride   CoCl2  6H2O     

.  .  s  .  . 

nitrate,  Co(NO3)2.6H2O  
sulphate,  CoSO4.7H2O  
sulphide  CoS 

..  d  .. 
.  .  s  .  . 
.  .  s  .  . 

...(0°)84     (91°)  340    ... 

43 

..26  39  83.. 

insoluble  in  water  or  dilute 
acids 

insoluble  ;  soluble  in  acids 

..71  79  108.. 
..82  150  275.. 

.  .  14  23  75  .. 

insoluble  in  water  or  acids 

very  soluble  ;  melts  35.5° 

.  2.0  4.7....  27.5.. 
300  

(0°)  3.5  (25°)  11.  4  ..  (70°)  64 

2 

....  5.0 
....  4.8 
....  1.2 

....  0.6 
....  0.9 

Copper: 
carbonate   CuCO3 

s  .  . 

chloride,  CuCl2.2H2O  
nitrate,  Cu(NO3)2.6H2O  

sulphate,  CuSO4.5H2O  
sulphide  CuS  

.s,d. 

]  melts 
(    38° 
.  s,  e  . 
.  .  s  .  . 

Hydrogen  : 
arsenic  acid   H3AsO4  -iH2O 

..  d  .. 

boric  acid,  H3BO3  
iodic  acid,  HIO3  
oxalic  acid,  H2C2O4.2H2O  

.  .  s  .  . 
.  d,  s. 
.  .  s  .  . 

APPENDIX 


181 


~  ,                                   Formula  of 
crystallized  salt. 

Behavior  when 
exposed  to 
atmosphere. 

SOLUBILITY  IN  WATER. 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°                25°             100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Hydrogen  : 
phosphoric  acid,  H3PO4  

Iron  : 
chloride  (ous),  FeCl2.4H2O  
(ic)    FeCl3  6H2O  .    ... 

..  d  .. 

..d  .. 
..  d  .. 

very  soluble  ;  melts  37° 
..(15°)  67    ....  (80°)  100.. 

very  soluble  ;  melts  31° 
insoluble  ;  soluble  in  acids 
(18°)  82 

2 

carbonate  (ous),  FeCO3        .... 

.  .  s  .  . 

nitrate  (ous),  Fe(NO3)2.6H2O  .  . 
(ic),  Fe(NO3)3.9H2O    .. 
sulphate  (ous),  FeSO4.7H2O  .  . 
(ic),  Fe2(SO4)3    

.  Ox  . 
..  d  .. 
e,  Ox 
..  d  .. 

very  soluble  ;  melts  47° 

(0°)  16..  (30°)  33.  ..(90°)  43 

very  soluble 
insoluble  ;  soluble  in  acids 

50  200.. 
..0.5  1.0  4.8  .. 
..'  0.0001  
07          11          33 

....  1.5 

....0.02 
..0.000003 
.    .  0  05 

.  .  s  .  . 

Lead: 
acetate,  Pb(C2H3O2)2.3H2O    .  . 
bromide,  PbBr2  
carbonate,  PbCO3  
chloride,  PbCU 

.  .  s  .  . 
.  .  s  .  . 
.  .  s  .  . 
.  .  s  .  . 

hydroxide,  Pb(OH)2 

001 

.  .  .00004 

iodide,  PbI2  
nitrate,  Pb(NO3)2  
sulphate,  PbSO4  

.  .  s  .  . 
.  .  s  .  . 
.  .  s  .  . 

.  .  0.04  .  .  .  0.08  0.44  . 
...38  59  132.. 
0004 

.  .  .  0.002 
....  1.4 
.  .  0.00013 

sulphide,  PbS  . 

.  .  s  .  . 

insoluble  ;  soluble  in  concen- 
trated strong  acids 

..  1.5  1.3  0.7  .. 

(13°)  5  5 

....0.17 
08 

Lithium  : 
carbonate,  Li2CO3  
bicarbonate,  LiHCO3    

.  .  s  .  . 
.  .  s  .  . 

chloride,  LiCl  
hydroxide,  LiOH.H2O   
nitrate,  LiNO3  3H2O    

..  d  .. 
CO2 
d 

...67  82  128.. 
..12.7  ...  12.9  ....17.5  . 
(0°)54  (30°)  138    (70°)  176 
...35  34  30  .. 

02 

....  13.3 
....  5.0 
....  7.3 

....  2.8 

0.01 

sulphate,  Li2SO4    

Magnesium  : 
carbonate,  MgCO3.3H2O  

.  .  s  .  . 
.  .  s  .  . 

chloride,  MgCl2.6H2O  
hydroxide   Mg(OH)2  

..  d  .. 
CO2 
..d  .. 

.  .  e  .  . 

.  .  s  .  . 

...53  57  73  .. 
0.001  
..(OP)  67  (40°)  85  .. 
...27  39  74  .. 

insoluble  ;  soluble  in  acids 

...63  77  115.. 

....  5.1 
...0.0002 
....  4.0 

....  2.8 

5 

nitrate,  Mg(NO3)2.6H2O  
sulphate,  MgSO4.7H2O    

Manganese  : 
carbonate,  MnCO3  

chloride,  MnCl2.4H2O  

.e,  d. 

182 


APPENDIX 


c  ,                             Formula  of 
crystallized  salt. 

Behavior  when 
exposed  to 
atmosphere. 

SOLUBILITY  IN  WATER. 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°                 25°            100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Manganese  : 
nitrate   Mn(NO3)2  6H2O    

..  d  .. 

(0°)102  (25°)  166    (35.5)331 
.  .  53  65  32  .. 

insoluble  ;  soluble  in  dilute 
acids 

0.0002  

5 
4 

.  .  0.00001 
....  0.2 
.  .  .00001 

sulphate,  MnSO4.5H2O    

.  s,  e  . 
.  Ox  . 

sulphide,  MnS  

Mercury  : 
chloride  (ous)   HgCl  

.  (ic)  HgCl2  

.  .  s  .  . 

..3.7  7.4  61  .. 
0  005  

nitrate  (ous),  HgNO3.H2O  
(ic),Hg(N03)24H20... 

.  s,  e  . 
..  d  .. 

}  very  soluble  in  a  little  water 
>  much  water  ppts.  basic  salt 
)  very  soluble  in  HNO3 
0.006  

...0.0001 

sulphide  (ic)  HgS 

s 

insoluble  ;  insoluble  in  con- 
centrated acids 

insoluble  ;  soluble  in  acids 
..54  67  88  .. 
(0°)80   .(20°)  96    (95°)  233 
(0°)  27  .  .  (30°)  43  ...  (99°)  77 

insoluble  in  water  or  dilute 
acids 

(6°)188  (14°)  230     (G2°)492 
..3.1  8.0  50  .. 
..54  68  104.. 
..  89  ....  113  156.. 
(0°)  22  .  .  (25°)  36  .  .  .(60°)  60 
..3.1  8.2  56  .. 
..28  36  57  .. 
..59  64  79  .. 
..  .5  16  89  .. 

4 
6 
2 

25 

Nickel  : 
carbonate   NiCO<> 

s 

chloride,  NiCl2.6H2O   
nitrate   NiNO3  6H2O 

.  s,  d  . 
.s,d. 
.  e  .  . 

sulphate,  NiSO4.7H2O  . 

sulphide,  NiS 

s 

Potassium  : 
acetate  KC2H3O2 

d 

bromate,  KBrO3     
bromide,  KBr  
carbonate,  K2CO3.1£H2O  
(bi)  carbonate,  KHCO3  

.  .  s  .  . 

.  .  s  .  . 

..d  .. 

.  .  s  .  . 

....0.38 
....  4.6 
....  5.9 
....  2.8 
....0.52 
....  3.9 
2.7 
....  0.4 

chlorate,  KC1O3  
chloride,  KC1    
chromate,  K2CrO4  

(bi)  rhromate   K2Cr2Oy 

.  .  s  .  . 
.  .  s  .  . 

.  .  s  .  . 

s 

fluoride   KF  2H2O 

d 

(18°)  92        .    .    . 

12.4 

hydroxide,  KOH.2H2O    
iodate,  KIO3  
iodide,  KI   

manganate    K2MnO4 

..  d  .. 
.  .  s  .  . 
.  .  s  .  . 
.  d  .. 

..  97  ....  119  178.. 
..4.7  9.9  32  .. 
..128....  148  208.. 
very  soluble 

..13  37  246.. 
.     38  

18 
....0.35 
....  6.0 

....  2.6 
1.6 

nitrate,  KNO3  
oxalate  K2C2O4  H2O 

.  .  s  .  . 

g 

perchlorate  KC1O4 

1.5  

0.11 

permanganate,  KMnO4  

.  .  s  .  . 

(0°)  2.8..  (25°)  8.0  ..(65°)  25 

....0.33 

APPENDIX 


c  i.                              Formula  of 
crystallized  salt. 

Behavior  when 
exposed  to 
atmosphere. 

SOLUBILITY  IN  WATER. 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°                 25°                 100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Potassium  : 
sulphate   K2SO4 

7             12             24 

062 

(bi)  sulphate,  KHSO4  
sulphide,  K2S.5H2O  
sulphite,  KoSO3.2H9O  

(o°)  36..  (20°)  51  .(100°)  122 

very  soluble 
very  soluble 

(0°)0.7.  (25°)  1.1..  (80°)  2.5 

0  00001  

....  3.5 

....0.06 
.  0  0000006 

..  d  .. 
..  d  .. 

Silver  : 
acetate,  AgC2H3O2   

.  .  s  .  . 

bromide,  AgBr 

g 

carbonate,  AgoCO^ 

s 

0.003  

...0.0001 
....  0.6 
.  .  0  00001 

chlorate,  AgClOs 

s 

15  

chloride,  AgCl    

s 

0.0002  

chromate,  Ag2CrC>4         .    . 

s 

0.002  

.  .  000015 

fluoride,  AgF  

d 

(16°)  182  

....  13.5 
.  .  0.00014 

iodate,  AglOs     ... 

S    . 

0.005  

iodide.  Agl 

0  0000003 

0  00000001 

nitrate,  AgNO3  
oxide,  AgoO,  dissolves  as  AgO  H 
perchlorate,  AgClO4 

.  .  s  .  . 
.  .  s  .  . 
d 

..122....  257  952.. 
0.0025  

very  soluble 

.  .  (18°)  073  ...  (ioo°)  15.. 

....  8.4 
...0.0002 

.  .  0  024 

sulphate   AffoSOa. 

s 

sulphide,  AgS    

s 

insoluble  in  water  or  acids 

(0°)  34..  (25°)  53...  (40°)  65 

(5°)  1.3  .(30°)  3.9..  (10(>°)  53 
..73  87  118.. 
..70  28  ..    .    46 

.  .  6 

Sodium  : 
acetate,  NaC2H3O2.3H2O  
(tetra)  borate 
(borax),  Na2B4O7.10H«,O  
bromide,  NaBr.2H2O    
carbonate,  Na2CO3.10H2O  
(bi)  carbonate,  NaHCO3  
chlorate,  NaClO3   
chloride,  NaCl    
chromate,  Na2CrO4.10H2O  
(bi)  chromate,  Na2Cr2O7.2H2O 
fluoride   NaF 

.  s,  e  . 

.  .  s  .  . 
.  .  s  .  . 

Q 

0.15 
....  0.9 
1  8 

.  .  S    .  . 

.  .  s  .  . 
.  .  s  .  . 
.  .  e  .  . 
..  d  .. 

g 

(0°)  6.9..  (25°)  10..  .(60°)  16 
..  82  105  233.. 
..36  36  40  .. 
(0°)  32..  (21°)  90  (ioo°)  126 
..(0°)16<5....    (98°)433  .. 
(2l°)4.2  

....  1.1 
....  6.4 
....  5.4 
....  3.3 
....  50 
1  l 

hydroxide    NaOH  H2O 

d 

.  .  42  114  348 

21 

iodide,  NaI.2H2O  
nitrate,  NaNO3  

.e,  d. 
.  .  s  .  . 

g 

..159....  184  302.. 
..73  92  178.. 
.  .  (15°)  3  2          (100°)  6  3 

....  8.1 
....  7.4 

024 

permanganate,  NaMnO4-3H2O 
sulphate,  Na2SO4.10H2O  
(bi)  sulphate,  NaHSO4.H2O  .. 

..  d  .. 
.  .  e  .  . 
..  d  .. 

very  soluble 

(0°)  5.0  (32.75°)  50.65  (100°)  43 
..(25°)29  (ioo°)  50  .. 

....  1.2 

1 84 


APPENDIX 


c  |.                              Formula  of 
crystallized  salt. 

Behavior  when 
exposed  to 
atmosphere. 

SOLUBILITY  IN  WATER. 

Grams  anhydrous  salt  per 
100  grams  water  in  a  sat- 
urated solution  at 
0°               25°               100° 

Mols  per  liter 
of  solution 
at  laboratory 
temperature. 

Sodium  : 
sulphide,  Na2S.9H2O    

(  Ox 
)d,e 
.  .  e  .  . 

s,  e 

.  .  s  .  . 

very  soluble 
(0°)  14  .  .  (20°)  27  ...  (40°)  50 

(10°)  60..  (25°)  76..  (45°)  124 
0  001      .    . 

2 
5 

sulphite,  Na2SO3.10H2O  
thiosulphate,  Na2S2O3.5H2O  .  . 

Strontium  : 
carbonate,  SrCOs  

..  0.00007 

chloride,  SrCl2.6H2O    
hydroxide,  Sr(OH)2.8H2O  
nitrate,  Sr(NO3)2.4H2O    
sulphate,  SrSO4  

'co2' 

.  .  e  .  . 
.  .  s  .  . 

..44  56  101.. 
..0.4  1.0  32  .. 
..40  79  101.. 
0  01  

....  3.0 
....006 
....  2.7 
...0.0006 

Tin: 
chloride  (ous),  SnCl2.2H2O  

Zinc: 
carbonate,  ZnCOs  
chloride,  ZnCl2.3H2O  
nitrate,  Zn(NO3)2.6H2O   
sulphate,  ZnSO4.7H2O  
sulphide,  ZnS   .    .  . 

f  s 
iOx 

.  .  s  .  . 
..  d  .. 
..d  .. 
.  .  e  .  . 

c 

(0°)  84           (15°)  270 

7 

..  0.0003? 
k.  .  .  .  9.2 
....  4.7 
....  3.1 

0.004  ?  
..208....  432  615.. 
95        .  ]  27   

..  42    .    .  .58  81  .. 

insoluble  ;  soluble  in  acids 

SUPPLEMENT 


In  the  preparation  of  the  third  edition  of  Synthetic  Inor- 
ganic Chemistry  it  has  seemed  desirable  to  proceed  slowly 
in  order  not  to  be  confronted  too  soon  after  the  completion 
of  the  edition  with  the  desirability  of  making  further  alter- 
ations and  additions. 

This  supplement  to  the  Second  Edition  is  therefore  printed 
that  the  Author  may  have  the  opportunity  to  test  new  ideas 
in  his  own  classes  and  that  other  users  of  the  book  may  have 
the  benefit  of  new  work  which  has  already  seemed  to  justify 
itself. 


187 


DIRECTIONS  FOR  WORK 

Preliminary  Report.  —  Before  beginning  work  on  a  prepa- 
ration the  student  should  have  a  clear  knowledge  of  the 
whole  procedure  and  should  understand  the  reactions  in- 
volved as  well  as  the  application  of  chemical  principles  to 
these  reactions. 

To  that  end  study  carefully  the  general  discussion  of  the 
preparation  as  well  as  the  procedure.  Then  write  in  the 
notebook  all  reactions,  and,  starting  with  the  given  amount 
of  the  principal  raw  material,  calculate  what  amounts  of  the 
other  substances  are  necessary  to  satisfy  the  equations. 
When  the  amount  specified  in  the  directions  is  different 
from  that  calculated,  state  the  reason  for  the  difference. 
Calculate  also  on  the  basis  of  the  equations  the  amount 
of  the  main  product  as  well  as  of  any  important  intermediate 
products  or  by-products. 

Present  this  preliminary  report  to  an  instructor  and  obtain 
his  approval  before  beginning  operations. 

Manipulation.  —  All  references  from  the  procedure  to  the 
general  notes  on  laboratory  manipulation  should  have  been 
studied  before  making  the  preliminary  report.  Indeed  the 
instructor  will  probably  make  sure  by  a  quiz  that  this  has 
been  done  before  he  accepts  the  preliminary  report. 

Laboratory  Record.  —  The  working  directions,  in  the  sec- 
tion entitled  procedure,  are  to  be  kept  at  hand  while  carrying 
out  the  manipulations.  These  directions  do  not  need  to 
be  copied  in  the  laboratory  notebook;  but  it  is  essential, 
nevertheless,  to  keep  a  laboratory  record  in  which  are  entered 
all  important  observations  and  data;  such  as,  for  example, 

189 


I QO  SUPPLEMENT 

appearance  of  solutions  (color,  turbidity);  appearance  of 
precipitates  or  crystals  (color,  size  of  grains,  crystalline 
form);  results  of  all  weighings  or  measurements;  number 
of  recrystallizations;  results  of  tests  for  purity  of  materials 
and  products,  etc. 

Questions.  —  The  section  under  this  title  gives  suggestions 
for  study,  which  involves:  (i)  laboratory  experiments  and 
direct  entries  in  the  laboratory  notebook;  (2)  consultation 
of  reference  books,  of  which  all  that  are  necessary  will  be 
found  upon  the  shelf  in  the  laboratory;  (3)  original  reasoning. 

The  answers  to  the  questions  should  be  written  in  the 
laboratory  notebook  following  the  entries  for  the  exercise, 
and  this  book  should  be  submitted  at  the  same  time  as  the 
preparation  for  the  approval  of  an  instructor. 

General  Questions.  —  Besides  the  specific  study  questions 
for  each  preparation  there  are,  accompanying  each  group 
of  exercises,  general  questions  relating  to  the  whole  group; 
and  these  are  to  be  worked  out  by  every  student.  The 
answers  to  these  questions  are  to  be  written  on  a  certain 
prescribed  kind  of  paper  and  handed  in  at  the  office,  neatly 
bound,  within  the  times  which  will  be  posted. 

Use  of  Time  in  Laboratory.  —  In  preparation  work  it  is 
frequently  necessary  to  wait  for  considerable  periods  of  time 
for  evaporations,  crystallizations,  etc.,  to  take  place.  This 
time  may  be  utilized  for  work  upon  the  study  questions 
and  experiments,  but  even  then  it  is  advisable  to  have  usu- 
ally more  than  a  single  preparation  under  way.  Thus  no 
time  need  be  wasted  by  the  energetic  student  who  plans 
his  work  well. 

Yield  of  Product.  —  Where  possible  the  methods  employed 
in  these  preparations  resemble  those  actually  used  on  an 
industrial  scale;  where  this  is,  however,  impossible  on  the 
limited  scale  of  the  laboratory,  mention  is  made  of  the  fact, 
with  reasons  therefor.  On  account  of  the  limitations  con- 


SUPPLEMENT  IQI 

nected  with  work  on  a  laboratory  scale,  it  is  of  course  im- 
possible to  get  as  high  percentage  yields  as  could  be  obtained 
on  a  large  commercial  scale.  The  amounts  obtained  of 
each  preparation  are  to  be  weighed  and  recorded,  but  the 
chief  stress  is  to  be  laid  upon  the  excellence  of  the  product 
rather  than  upon  its  quantity. 


NOTES   ON    LABORATORY    MANIPU- 
LATION 

12.  AUTOMATIC   GAS   GENERATOR 

FOR   HYDROGEN,   HYDROGEN    SULPHIDE,   OR    CARBON    DIOXIDE. 

Supplies:   i  300-00.  generator  bottle  (thick  walled). 

i  calcium  chloride  drying  tube  (10  inches  long 
exclusive  of  stem;    i  inch  internal  diameter), 
i  reservoir  of  300  cc.  capacity. 
i  2-hole  rubber  stopper  to  fit  generator  bottle, 
i  i -hole  rubber  stopper  to  fit  drying  tube, 
i  foot  rubber  tube  to  connect  generator  bottle 

and  reservoir, 
glass  wool. 

This  apparatus  is  based  on  the  principle  of  the  familiar 
Kipp  generator  and  it  is  especially  suited  to  cases  in  which 
a  solution  is  to  be  saturated  with  the  gas  in  question,  as,  for 
example,  when  an  ammoniacal  solution  of  common  salt  is 
to  be  saturated  with  carbon  dioxide  in  the  preparation  of 
sodium  bicarbonate  by  the  Solvay  process. 

Directions  for  Setting  up  and  Starting  the  Generator.  —  As- 
semble the  apparatus  as  shown  in  the  diagram.  The  stem 
of  the  generator  tube  E  should  reach  flush  with  the  bottom 
of  the  stopper  but  not  below.  The  delivery  tube  C  should 
reach  nearly  to  the  bottom  of  the  generator  bottle  D.  Place 
the  requisite  amount  of  calcium  carbonate  (or  zinc,  or  ferrous 
sulphide)  in  the  generator  tube.  Then  insert  a  loose  plug 
of  glass  wool  F  about  i  J  inches  long  so  that  it  will  stand  about 

192 


SUPPLEMENT 


193 


midway  between  the  top  of  the  solid  material  and  the  stopper 
in  the  mouth  of  the  tube,  and  act  as  a  gas  filter  (to  remove 
acid  spray).  Pour  the  requisite  amount  of  acid  into  the 


AUTOMATIC  GENERATOR 


reservoir  A ;  clamp  the  reservoir  at  just  the  same  height  as 
the  generator  tube  and  pour  in  water  cautiously  until  the 
acid  rises  and  barely  touches  the  solid  in  the  generator  tube. 


IQ4  SUPPLEMENT 

The  generation  of  gas  will  now  begin  and  proceed  auto- 
matically. Do  not  pour  any  more  water  into  the  reservoir 
until  the  air  is  swept  from  the  receiving  flask  and  the  mouth 
of  the  latter  is  closed  tight  (see  below).  Then  add  not  more 
than  5  cc.  of  water.  Later  after  the  vigor  of  the  absorption 
has  slackened  the  reservoir  may  be  raised  to  a  higher  level 
to  give  a  greater  pressure  and  some  more  water  may  be 
added. 

If  the  solid  charge  or  the  acid  becomes  exhausted,  lower 
the  reservoir  to  a  little  below  the  generator  tube;  remove  the 
stopper  from  the  generator  tube  and  introduce  more  of 
the  solid  if  necessary;  if  the  acid  needs  renewal,  unclamp 
the  reservoir,  lower  it  and  invert  it  to  let  the  spent  liquor 
siphon  out  of  the  generator  bottle,  and  refill  with  acid  as 
in  the  first  charging. 

Directions  for  Using  the  Generator.  —  Place  the  solution 
to  be  saturated  in  a  flask  (G  in  diagram)  fitted  with  a  i-hole 
rubber  stopper  through  which  passes  a  delivery  tube  reaching 
to  the  bottom  of  the  flask.  Start  the  generator  in  action 
(see  directions  above).  Loosen  the  stopper  in  the  flask  a 
little  so  that  gas  escapes  until  all  the  air  originally  in  the 
generator  and  receiving  flask  is  swept  out.  Then  make 
the  stopper  tight;  the  gas  will  now  pass  in  as  rapidly  as  it 
can  be  absorbed  by  the  solution.  Shaking  the  receiving 
flask  will  greatly  increase  the  rapidity  of  absorption,  but 
observe  this  caution: 

At  the  outset  if  the  gas  is  drawn  too  rapidly  the  liquid  may 
rise  so  far  in  the  generating  tube  E  as  to  produce  too  violent  an 
action,  which  will  either  blow  out  the  stoppers,  or  cause  froth  to 
pass  over  into  the  receiver.  Therefore  be  cautious  not  to  sweep 
out  the  air  too  rapidly,  and  after  the  stopper  is  placed  firmly 
in  the  receiving  flask  be  cautious  during  the  first  75  minutes 
not  to  shake  the  receiver  too  strongly. 


SUPPLEMENT  1 95 

I.  POTASSIUM  NITRATE  FROM   SODIUM 
NITRATE  AND   POTASSIUM   CHLORIDE 

Read  the  discussion  on  pages  25-26  of  the  second  edition. 

In  the  following  procedure  equi-molal  amounts  of  sodium 
nitrate  and  potassium  chloride  are  taken  and  enough  water 
added  to  dissolve  at  the  boiling  temperature  all  of  the  sodium 
nitrate  taken  or  all  of  the  potassium  nitrate  which  could 
result  from  metathesis,  but  not  enough  to  dissolve  either 
the  potassium  chloride  taken  nor  the  sodium  chloride  wrhich 
could  be  formed  by  metathesis.  Nevertheless  after  this 
mixture  is  boiled  a  short  time  all  of  the  solid  potassium 
chloride  disappears  and  the  only  solid  salt  left  is  sodium 
chloride. 

The  mechanism  of  the  chemical  process  may  be  more 
easily  appreciated  if  it  is  represented  on  paper  in  the  fol- 
lowing fashion: 

KC1       -»KC1      <=±K+        CF 
NaNO3->NaNO3<=±NO3~    Na+ 

IT         IT 

KN03      NaCl 

IT 

NaCl 

The  formulas  printed  in  bold  face  type  stand  for  the  sub- 
stances in  the  solid  state,  —  those  in  common  type  for  sub- 
stances in  the-  dissolved  state.  Single  arrows  indicate  that 
the  reaction  runs  to  completion  in  that  direction  under  the 
conditions  prevailing  (the  boiling  temperature  is  supposed 
to  be  prevailing  in  the  above  representation).  The  double 
arrows  indicate  that  an  equilibrium  is  reached  and  no  sub- 
stance shown  on  either  side  of  the  arrows  disappears  from 
the  sphere  of  action.  If  the  conditions  were  to  be  shown  at 
o°  solid  potassium  nitrate  would  have  to  be  indicated  in 
equilibrium  with  the  dissolved  salt. 


196  SUPPLEMENT 

Materials:    crude  Chili  saltpeter  NaNOs,  100  grams. 

crude  potassium  chloride  KC1,  88  grams. 
Reagent:       i%  AgNO3  solution. 
Apparatus:  35o-cc.  casserole. 

watch  glass. 

5-inch  funnel. 

perforated  filter  plate. 

8oo-cc.  suction  bottle  and  pump. 

platinum  wire. 

Procedure.  —  Place  100  grams  NaNO3  and  88  grams  KC1 
in  a  350-cc.  casserole.  Add  125  cc.  of  water,  cover  with  a 
watch  glass,  and  place  over  a  low  flame.  Keeping  an  eye  on 
the  casserole  to  see  that  the  contents  do  not  boil,  prepare  a 
suction  filter  according  to  Note  4  (b),  on  page  7.  Then  raise 
the  flame  under  the  casserole  and  watch  it  until  boiling  com- 
mences. Lower  the  flame  and  let  the  mixture  boil  gently 
just  one  minute,  keeping  the  watch  glass  over  the  casserole 
to  prevent  too  much  evaporation  of  water.  While  it  is  at 
the  boiling  temperature,  pour  (see  Figure  i,  page  6)  the 
mixture  from  the  casserole  onto  the  suction  filter  after  first 
starting  a  gentle  suction.  Quickly  scrape  most  of  the  damp 
salt  onto  the  filter  and  suck  out  as  much  of  the  liquid  as 
possible.  Then  return  the  solid  salt,  which  is  mostly  NaCl, 
to  the  casserole.  Pour  the  solution  into  a  beaker  and  cool 
it  to  15°  or  below  by  setting  it  in  a  pan  of  cold  water  or  snowr. 
Separate  the  crystals  of  KNOs  from  the  cold  liquor  by  means 
of  the  suction  filter,  observing  last  sentence  of  Note  3  on  page 
7,  and  pour  the  liquor  into  the  casserole  containing  the  first 
crop  of  NaCl  crystals.  Bring  the  solution  to  boiling  point 
and  boil  gently  three  minutes  without  a  watch  glass  over  the 
casserole,  thus  allowing  some  of  the  water  to  escape  by 
evaporation.  Then  filter  at  the  boiling  temperature  exactly 
as  in  the  first  instance.  Cool  the  filtrate  and  collect  a  second 
crop  of  KNOa  crystals,  adding  them  to  the  first  crop  and 


SUPPLEMENT  197 

pouring  the  liquor  into  a  flask  labelled  "Mother  Liquors." 
Examine  the  two  kinds  of  crystals,  tasting  them  and  using 
a  microscope.  Draw  pictures  in  the  notebook  of  the  crys- 
tals as  seen  in  the  microscope.  Dissolve  about  o.i  gram 
of  the  supposed  KN03  in  2  cc.  of  water  and  test  for  chloride 
by  adding  i  drop  of  AgNOs  solution.  Considerable  chloride 
will  be  found  and  the  product  must  be  purified  by  recrys- 
tallization.  Weigh  the  crystals  roughly  while  they  are  still 
moist,  add  one  half  their  weight  of  hot  water,  and  warm  until 
solution  is  complete.  Cool  to  below  15°  and  separate  the 
crystals  from  the  mother  liquor,  adding  the  latter  to  the 
reserve  flask.  Test  as  above  to  see  if  this  crop  of  crystals 
is  free  from  chloride.  If  not  repeat  the  recrystallization  as 
many  times  as  is  necessary  to  get  a  perfectly  pure  product. 
A  little  of  this  when  dissolved  should  give  no  turbidity  with 
silver  nitrate  solution,  and  when  held  in  the  flame  on  a  plati- 
num wire  should  color  it  the  violet  color  characteristic  of 
potassium  with  none  of  the  yellow  sodium  color.  Spread  the 
preparation  on  filter  paper  or  an  unglazed  porcelain  plate 
and  allow  it  to  dry  by  standing  exposed  to  the  air;  then  put 
up  the  salt  in  a  test  tube  or  a  small  bottle  and  label  it  neatly. 
If  the  final  yield  of  pure  product  is  not  satisfactory  in  amount 
the  collected  mother  liquors  should  be  boiled  down  to  about 
100  cc.,  and  used  as  the  starting  point  in  a  repetition  of  the 
above  procedure.  30  grams  may  be  regarded  as  a  very 
satisfactory  yield. 

The  sequence  of  the  operations  in  this  preparation  can  be 
followed  rather  more  readily  in  the  tabulated  procedure 
shown  on  the  following  page. 


198 


SUPPLEMENT 


TABULATED   PROCEDURE 

Treat  100  grams  NaNO3  and  88  grams  KC1  with  125  cc.  water; 
heat  to  boiling  and  boil  one  minute;  filter  hot.     Do  not  rinse  out 
dish  but  keep  it  for  second  boiling. 

On  filter: 
NaCl,  dirt, 
some  KNO3: 
transfer  back 
to  dish  in 
which  first 
boiling  was 
made 
(i) 

Filtrate:  cool  and  filter 

Crystals: 
impure 
KN03 

(2) 

Filtrate  is  saturated  with  KNO3  and 
NaCl.     Pour  into  dish  in  which  origi- 
nal mixture  was  boiled  and  to  which 
impure  NaCl  (i)  was  added.     Bring  to 
boil,  boil  3  minutes,  and  filter  hot. 

On  filter:      Filtrate: 
NaCl 

cool  and  filter 

and  dirt      Crystals: 
fairly  free    impure 
from          KNQ3 
KNO,            (4) 

Filtrate  is  satu- 
rated with  KNO3 
and  NaCl. 
Save  temporarily 
in  flask  labelled 
"Mother  Liquors" 
(5) 

RECRYSTALLIZATION 

Unite  impure  KNO3  (2)  and  (4)  ;  heat  with  one  half  their  weight 
of  water  until  dissolved;  cool  and  filter. 

Crystals:  Nearly  pure  KNO3. 
Recrystallize    repeatedly  until 
entirely  pure,  adding  all  mother 
liquors   to   (5)    in   the   reserve 
flask. 

Filtrate  contains  nearly  all  of 
the  NaCl  from  the  impure  prod- 
uct, and  is  saturated  with  KNO3; 
add  to  (5)  in  the  reserve  flask. 

Discard  mother  liquors  (5)  if 
factory. 

the  yield  of  pure  KNO3  is  satis- 

Questions 

1.  Define  metathesis. 

2.  When  a  metathetical  reaction  is  carried  out  in  the 
wet  way,  why  is  the  solubility  of  the  substances  involved 
of  importance?     Explain  why,  according  to   this  point  of 
view,    the   reactions   AgNO3  +  KC1  =  AgCl  +  KNO3    and 
BaCl2  +  Na2SO4  =  BaSO4  +  2NaCl   are   much   more   com- 
plete than   the  reaction  NaNO3  +  KC1  =  KN03  +  NaCl. 


SUPPLEMENT  199 

3.  Explain    why    fewer    operations    should    be    required 
to  prepare  potassium  nitrate  from  potassium  sulphate  and 
barium  nitrate  than  by  the  foregoing  procedure. 

4.  Explain  why  all  of  the  solid  salt  should  change  to 
NaCl  when  the  original  materials  are  boiled  with  insufficient 
water  to  dissolve  all  of  either  the  NaCl  or  the  KC1. 

5.  In  the  tabulated  procedure  what  is  the  advantage  of 
adding  the  impure  NaCl  (i)  to  the  second  mother  liquor 
instead  of  discarding  it? 

3.   SODIUM  BICARBONATE  BY  THE  AMMONIA 
(SOLVAY)   PROCESS. 

Read  the  discussion  on  pages  31,  32  of  the  Second  Edition. 
Materials:  table  salt,  59  grams. 

i4-normal  ammonium  hydroxide,  71  cc. 
cracked  marble,  105  grams. 
i2-normal  hydrochloric  acid,  175  cc. 
Apparatus:  automatic  gas  generator  (see  Note  12  on  p.  192 

of  Supplement). 

suction  filter  (see  Note  4,  a  and  b  on  p.  7). 
75O-CC.  flask  equipped  with  i-hole  rubber  stop- 
per, delivery  tube  reaching  to  bottom,  and 
1 8  inches  of  rubber  delivery  tube. 
3oo-cc.  flask  with  stopper. 

Procedure.  —  Place  the  salt,  the  ammonium  hydroxide, 
and  130  cc.  of  water  in  the  smaller  flask  and  shake  vigor- 
ously until  the  salt  is  dissolved.  Pour  the  solution  through 
a  filter  into  the  750-0:.  flask  (large  plaited  filter  for  speed). 
Use  this  flask  as  the  absorption  vessel  and  connect  it  with 
the  carbon  dioxide  generator.  Charge  the  generator  with 
the  cracked  marble  and  hydrochloric  acid  and  proceed  to 
saturate  the  solution  with  carbon  dioxide  following  with 
care  the  directions  for  starting  and  using  the  generator 
(Note  12).  After  10  minutes  commence  shaking  the  flask 


2OO  SUPPLEMENT 

very  cautiously  and  after  5  minutes  more,  if  it  is  found  to  be 
safe,  shake  vigorously.  Let  the  absorption  continue  until 
practically  no  more  gas  passes  into  the  absorption  flask  even 
with  vigorous  shaking.  If  the  shaking  has  been  continuous, 
this  point  will  be  reached  within  one  hour.  It  is  equally 
as  well  to  allow  the  absorption  to  proceed  by  itself  over  night 
and  to  shake  next  day  to  complete  the  saturation  with  the 
gas.  When  the  absorption  is  complete  collect  the  precipi- 
tated sodium  bicarbonate  on  the  suction  filter  (Notes  3,  4a, 
and  4b),  drain  it  thoroughly  with  suction,  stop  the  suction, 
pour  over  the  surface  of  the  product  15  cc.  of  cold  water,  and 
after  this  has  soaked  in  apply  the  suction  again.  Wash 
a  second  time  with  15  cc.  of  cold  water  exactly  as  at  first. 
Spread  the  drained  product  on  a  clean  unglazed  porcelain 
plate  (or  on  filter  paper  spread  on  a  folded  newspaper)  and 
leave  it  24  hours  to  dry.  Test  the  preparation  for  chloride 
by  dissolving  about  o.i  gram  in  a  little  water,  acidulating 
slightly  with  nitric  acid,  and  adding  a  drop  of  silver  nitrate 
solution.  There  will  be  considerable  clouding. 

Questions 

1.  What  is  the  purpose  of  washing  the  product  with  water? 
How  much   sodium   bicarbonate   is  lost  in  this   way   (see 
solubility  table)? 

2.  Why  must  the  solution  be  acidulated  with  nitric  acid 
before  testing  with  silver  nitrate? 

3.  Why  does  shaking  greatly  increase  the  rate  of  absorp- 
tion? 

4.  How  do  you  explain  the  heat  produced  in  the  absorption 
flask? 

5.  How  can  you  prepare  sodium  carbonate  from  sodium 
bicarbonate? 

6.  Why  cannot  potassium  bicarbonate  be  effectively  pre- 
pared from  potassium  chloride  by  the  ammonia  process? 


SUPPLEMENT  201 

(Look  up  the  solubility  of  potassium  bicarbonate.)  What 
process  may  be  used  to  prepare  potassium  carbonate  from 
this  source? 

7.  What  is  an  acid  salt?    How  does  a  solution  of  an  acid 
salt  such  as  KHSO4  behave  toward  litmus?    Test  the  be- 
havior of  solutions  of  NaHCO3  and  Na^COa  toward  litmus. 
Explain  the  cause  of  this  behavior. 

8.  Would  a  precipitate  of   sodium   bicarbonate   form  if 
carbon  dioxide  were  passed  into  a  solution  of  sodium  chloride 
alone?    Explain  the  part  played  by  the  ammonia  in  the 
formation  of  the  product. 

9.  Explain  how  a  given  amount  of  ammonia  may  be  used 
over  and  over  again. 

3-A.   SODIUM   CARBONATE  FROM   SODIUM 
BICARBONATE 

Heat  the  sodium  bicarbonate  obtained  in  No.  3  until  it 
is  converted  into  sodium  carbonate.  Compare  the  weight 
obtained  with  that  calculated. 

3-B.  CAUSTIC  SODA  FROM  SODIUM  CARBONATE 

Apparatus:    8-inch  dish, 
suction  filter, 
burette  with  normal  HC1. 
i5-cc.  pipette. 
500-cc.  bottle  with  rubber  stopper. 

Read  the  discussion  and  procedure  of  No.  2,  Caustic 
Potash  from  Wood  Ashes,  and  adapt  that  method  to  the 
conversion  into  sodium  hydroxide  of  the  sodium  carbonate 
obtained  in  No.  3~a.  Keep  the  sodium  hydroxide  solution 
in  the  rubber  stoppered  bottle. 

Answer  Questions  i  and  2  under  No.  2. 


202  SUPPLEMENT 

5.   AMMONIUM  BROMIDE 

The  solution  contains  practically  pure  ammonium  bromide 
all  of  which  should  be  recovered.  Evaporate  to  complete 
dryness  and  pulverize  the  caked  salt  thereby  obtained.  If 
preferred  the  first  part  of  the  evaporation  may  be  hastened 
by  heating  to  gentle  boiling  with  a  small  flame  held  constantly 
in  the  hand.  It  is  unsafe  to  put  a  flame  under  the  dish  and 
leave  it,  for  the  solution  may  "  bump  "  and  spatter.  Further- 
more as  soon  as  the  salt  is  dry  it  will  volatilize  freely  itself. 

5-A.   MAGNESIUM  NITRIDE  AND  AMMONIUM 
SALT  FROM  ATMOSPHERIC  NITROGEN 

Active  metals,  as  those  of  the  alkali  and  alkaline  earth 
families,  when  heated,  combine  readily  with  nitrogen.  In  this 
preparation  powdered  magnesium  in  a  closely  packed  mass 
is  allowed  to  react  with  air.  The  oxygen  reacts  with  the 
upper  layers  and  only  nitrogen  penetrates  to  the  interior  of 
the  mass.  Thus  a  large  part  of  the  magnesium  should  be 
converted  to  nitride.  On  treatment  with  water  the  substance 
hydrolyzes  and  the  ammonia  given  off  can  be  absorbed  in  an 
acid  to  yield  ammonium  salt. 

The  mechanism  of  the  hydrolysis  of  magnesium  nitride 
is  probably  similar  to  that  of  such  salts  as  sodium  carbonate 
and  ferric  chloride,  and  it  would  therefore  appear  as  follows: 


Mg3N2->3Mg++ 

6H2O   ->6OH~  6H+ 

I  I 

3Mg(OH)2       2NH3 

If  magnesium  nitride  can  ionize  at  all  the  ions  which  it 
would  yield  are  obviously  Mg++  and  N  as  shown  in  the 
upper  horizontal  equation.  The  fact  that  such  an  ion  as 
N  is  entirely  unfamiliar  does  not  weaken  our  belief  in 


SUPPLEMENT  203 

the  above  mechanism,  because,  as  the  direction  of  the  arrow 
shows,  the  N  ion  is  formed  only  to  be  removed  according 
to  the  right-hand  vertical  equation. 

Materials:     powdered  magnesium,  10  grams. 

dry  sand,  50  grams. 
Apparatus:    iron  crucible  and  cover  of  25  cc.  capacity. 

300-cc.  r.b.  flask. 

5<D-cc.  distilling  flask. 

filter  bottle  (or  500-0:.  flask). 

U-tube  |  inches  wide,  ;J  inches  tall. 

2  thistle  tubes. 

1  pinch  cock. 

2  2-hole  rubber  stoppers. 
i  i -hole  rubber  stopper. 

See  diagram  for  fittings. 

Procedure.  —  Weigh  an  iron  crucible  of  about  25  cc.  volume 
together  with  the  cover.  Pack  it  even  full  with  powdered 
magnesium,  tapping  the  crucible  on  the  desk  to  make  the 
powder  settle.  Weigh  the  filled  crucible;  it  should  hold 
10  grams  of  the  powder.  Place  the  cover  tightly  on  top, 
surround  the  crucible  with  a  cylinder  of  asbestos  i\  to  3 
inches  in  diameter  so  as  to  diminish  the  loss  of  heat  by  ra- 
diation. Heat  the  crucible  as  hot  as  possible  with  a  Tirrell 
burner  for  45  minutes.  After  it  cools  empty  the  crucible 
onto  a  piece  of  paper  and  note  the  white  MgO  on  top  and 
the  yellow  MgsN2  beneath.  Place  the  material  in  a  mortar, 
break  up  the  lumps,  add  25  grams  of  dry  coarse  sand,  and 
mix  well.  Then  place  25  grams  of  sand  in  the  bottom  of  a 
dry  3oo-cc.  flask  and  pour  the  mixture  from  the  mortar  on 
top  of  it.  Use  this  as  the  generating  flask  A  in  the  diagram. 
Have  the  rest  of  the  apparatus  and  connections  fitted  as 
in  the  diagram.  Pour  50  cc.  of  water  into  the  absorption 
flask  E.  The  water  should  seal  the  bottom  of  the  thistle 


2O4 


SUPPLEMENT 


APPARATUS  FOR  DISTILLATION  OF  AMMONIA  FROM  MAGNESIUM  NITRIDE 
A  =  Generating  flask.  B  =  Trap  to  catch  solid  matter  entrained  with 
gas  and  steam.  C  =  Connecting  tube  of  f-inch  internal  diameter. 
D  =  Hole  for  return  of  solid  and  condensed  water  to  flask.  E  =  Absorp- 
tion flask  with  pure  water  in  bottom.  F  =  Absorption  tube,  bend  sealed 
with  dilute  acid;  glass  beads  are  an  advantage. 


SUPPLEMENT  2O5 

tube  h  but  should  stand  about  J  inch  below  the  end  of  the 
delivery  tube  g.  Place  10  cc.  of  6-normal  HC1  in  the  ab- 
sorption tube  F  and  then  add  enough  water  to  seal  the  bend. 
Remove  the  stopper  and  fittings  from  the  generating  flask  A . 
Pour  water  into  the  thistle  tube  i  and  open  the  pinch  cock 
until  the  stem  of  the  tube  has  filled  with  water.  Replace 
the  stopper  in  the  flask  and  open  the  pinch  cock  to  admit  a 
single  drop  of  water.  Add  another  drop  as  soon  as  the 
reaction  subsides  and  continue  to  add  a  single  drop  at  a  time 
until  the  reaction  becomes  less  violent.  Finally  add  enough 
water  to  make  70  cc.,  rock  the  flask  until  the  contents  are 
thoroughly  mixed,  then  while  still  rocking  the  flask  apply 
a  small  flame  until  the  liquid  boils.  Boil  gently  for  15 
minutes.  Pour  together  the  contents  of  the  absorption 
flask  and  the  absorption  tube,  and,  using  litmus  as  an  indi- 
cator, add  enough  more  6-normal  HC1  to  just  neutralize  the 
ammonia.  Evaporate  the  solution  to  obtain  solid  ammonium 
chloride,  following  the  directions  for  Ammonium  Bromide 
on  page  202  of  this  Supplement. 

Questions 

1.  Experiment.  —  Burn  a  little  calcium  in  the  air  and  test 
the  ash  for  nitride.     How? 

2.  Give  reasons  for  regarding  the  action  of  magnesium 
nitride  with  water  as  an  example  of  hydrolysis.     Remember 
that  hydrolysis  is  the  exact  reverse  of  neutralization  and 
produces  an  acid  and  a  base  from  a  salt  and  water.     What  is 
the  acid  and  what  is  the  base  in  this  case? 

3.  Why  is  it  necessary  to  mix  the  magnesium  nitride  with 
an  inert  material  such  as  sand  before  adding  water? 

4.  The  layer  from  the  top  of  the  crucible  will  often  contain 
a  black  substance  as  well  as  a  white,  and  particularly  so  if 
the  gases  from  the  flame  entered  under  the  lid  of  the  crucible. 
What  is  this  black  substance  and  why  should  it  have  formed? 


206  SUPPLEMENT 

5.  On  the  basis  that  air  contains  4  volumes  of  nitrogen  to 
i  volume  of  oxygen  figure  what  fraction  of  the  magnesium 
would  be  converted  to  nitride  under  the  most  favorable 
conditions. 

6  AND  8.     TO  DRY  THE  PREPARATIONS  OF 

STRONTIUM   HYDROXIDE  AND   BARIUM 

HYDROXIDE 

Spread  the  moist  crystals  in  a  layer  J  inch  deep  on  a  large 
sheet  of  filter  paper.  Press  with  another  filter  paper,  then 
fold  the  edges  of  the  lower  paper  up  tightly  over  the  crystals 
and  set  the  whole  away  in  the  cupboard  (never  in  a  heated 
place)  until  the  next  exercise.  The  moisture  is  drawn  through 
the  paper  by  capillarity  and  evaporates  from  the  outside, 
while  the  paper  to  a  large  extent  prevents  the  carbon  dioxide 
from  reaching  the  preparation. 

9.   BORIC  ACID 

Procedure.  —  Dissolve  100  grams  of  borax  Na2B407.ioH2O 
in  300  cc.  of  boiling  water.  Add  a  few  drops  of  methyl 
orange  solution  (a  dye  which  is  yellow  in  neutral  or  alkaline 
solution  but  is  pink  in  distinctly  acid  solution)  and  add  i2-n 
HC1  until  the  color  of  the  dye  has  changed  through  an  orange 
to  a  distinct  pink  and  addition  of  i  cc.  more  of  the  acid  does 
not  increase  the  pink  tone.  Let  the  solution  cool  to  15°  or 
below,  drain  the  crystals  on  a  suction  filter.  If  the  filtrate  is 
not  distinctly  pink  (showing  it  to  be  acid)  add  enough  i2-nHCl 
to  make  it  so,  shake  vigorously,  and  add  any  crystals  thus 
obtained  to  the  main  crop.  Dissolve  the  crystals  in  300  cc. 
of  boiling  water,  filter  if  not  clear,  crystallize  by  slow  cooling 
with  occasional  stirring  if  crystals  cake  together  too  much. 
Collect  the  crystals  and  let  them  dry  at  room  temperature. 

Note.  —  The  above  procedure  yields  fine  granular  crystals. 
Lustrous  flaky  crystals  can  be  obtained  if  a  little  grease  is 


SUPPLEMENT  207 

present  in  the  crystallizing  solution.  To  this  end  dissolve 
i  gram  of  soap  shavings  in  300  cc.  of  boiling  water  and  use 
this  solution  to  dissolve  the  boric  acid  for  the  recrystallization. 
This  procedure  is  not  recommended  if  purity  of  product  is 
required  rather  than  an  attractive  appearance. 

II-A.    ALUMINUM  NITRIDE  AND  AMMONIA 
FROM  ATMOSPHERIC  NITROGEN 

In  a  recently  developed  process  named  after  the  inventor, 
Serpek,  aluminum  oxide  and  carbon  are  heated  in  an  electric 
furnace  in  an  atmosphere  of  nitrogen  (or  rather  of  producer 
gas,  which  contains  nitrogen  and  carbon  monoxide).  The 
reaction  is  essentially  as  follows: 

A1203  +  N2  +  3  C  ->  2  A1N  +  3  CO. 

Now  it  is  well  known  that  aluminum  oxide  cannot  be  re- 
duced to  metal  by  means  of  carbon  but  that,  at  very  high 
temperature  the  oxygen  may  be  withdrawn  and  carbon  sub- 
stituted for  it,  thus  yielding  the  carbide: 

2  A1203  +  9  C  ->  A14C3  +  6  CO. 

In  an  atmosphere  of  nitrogen,  however,  the  place  of  that 
part  of  the  carbon  that  unites  with  the  aluminum  may  be 
taken  by  the  nitrogen,  and  aluminum  nitride  is  thus  obtained. 

The  manipulation  of  a  powerful  enough  electric  furnace 
on  a  small  scale  for  a  laboratory  preparation  is  scarcely  feasi- 
ble. But  if  we  start  with  metallic  aluminum  instead  of  the 
oxide  we  may  easily  effect  its  combination  with  nitrogen. 

Materials:  finely  powdered  aluminum,  16  grams.  The  ma- 
terial sold  for  use  as  pigment  and  often  labelled 
Aluminum  Bronze,  although  it  is  nearly  pure 
aluminum,  is  suitable  for  the  purpose.  The  oil 
which  still  adheres  from  the  grinding  is  of  no 
disadvantage. 


208  SUPPLEMENT 

Materials:    lamp  black  f  gram. 

sodium  hydroxide,  30  grams.  Use  material 
which  comes  in  sticks  or  large  lumps.  It 
should  contain  at  least  95%  of  NaOH  and 
not  more  than  a  trace  of  carbonate.  Avoid 
the  granulated  material  for  this  preparation. 
Apparatus:  iron  crucible  and  cover  of  25  cc.  capacity. 

dark-colored  goggles  (recommended  to  protect 
the  eyes  from  the  blinding  light). 

looo-cc.  round  bottom  flask. 

50-cc.  distilling  flask. 

filter  bottle  (or  500-0:.  flask). 

U-tube  J  inches  wide,  yj  inches  tall. 

i  thistle  tube. 

1  2-hole  rubber  stopper. 

2  i -hole  rubber  stoppers. 

See  diagram  under  No.  5~a  for  assembling  the  apparatus. 
A  larger  generating  flask  A  is  used.  The  stem  of  the  thistle 
tube  i  reaches  to  bottom  and  the  stop  cock  is  dispensed  with. 
The  safety  tube  h  in  absorption  flask  is  unnecessary. 

Procedure.  —  Mix  the  16  grams  of  aluminum  and  f  gram 
of  lamp  black  thoroughly  by  grinding  together  in  a  mortar. 
Place  the  mixture  in  the  iron  crucible,  packing  it  down  as 
compactly  as  possible  by  tapping  the  crucible  on  the  desk 
top.  The  joint  of  the  cover  should  be  fairly  tight  to  exclude 
air  during  the  preliminary  heating.  If  a  little  moist  shredded 
asbestos  is  packed  around  under  the  edge  of  the  cover  before 
the  latter  is  pressed  down  onto  the  crucible,  such  a  tight 
joint  is  obtained.  Heat  the  crucible  with  a  Tirrell  burner 
rather  cautiously  at  first  until  the  oil  upon  the  aluminum 
has  ceased  to  give  off  inflammable  gas.  Then  set  the  Tirrell 
burner  under  the  crucible  and  start  heating  the  top  of  cover 
with  a  blast  lamp.  When  the  cover  is  bright  red  remove 


SUPPLEMENT  2OQ 

the  lamp  from  underneath;  keep  heating  the  cover  as  strongly 
as  possible  until  the  under  part  of  crucible  has  cooled  so  that 
it  is  no  longer  visibly  red.  Remove  the  cover  quickly  and 
play  the  tip  of  blast  flame  upon  one  part  of  the  surface  of 
the  charge  until  it  ignites  and  becomes  blindingly  incan- 
descent. The  reaction  is  then  self-sustaining  and  will  spread 
gradually  to  every  part  of  the  crucible. 

After  examining  the  cooled  product  in  the  crucible  pul- 
verize it  and  transfer  it  to  the  generating  flask.  Place  75  cc. 
of  water  in  the  absorption  flask  and  30  cc.  of  6-normal 
H2SO4  in  the  absorption  tube.  Add  50  cc.  of  water  to  the 
generating  flask.  Make  sure  that  this  flask  is  supported  in 
such  a  way  that  a  pan  of  cold  water  can  be  raised  from  under- 
neath to  cool  the  flask  whenever  the  reaction  becomes 
violent  and  has  to  be  checked.  Dissolve  30  g.  of  NaOH  in 
50  cc.  of  water  and  add  10  cc.  of  this  solution  to  the  generat- 
ing flask.  Warm  this  flask  until  a  vigorous  reaction  begins. 
Remove  the  flame  and  add  the  rest  of  the  NaOH  little  by 
little.  Thereupon  warm  the  flask  during  J  hour  just  enough 
to  keep  the  contents  boiling  gently.  It  is  now  well  to  let 
the  flask  cool  a  little  and  to  leave  it  stoppered  over  night. 
It  may  then  be  reconnected  and  boiled  for  15  minutes.  If 
this  is  not  feasible  the  boiling  should  be  continued  imme- 
diately for  another  |  hour.  Save  the  solution  remaining  in 
the  flask  as  the  starting  material  for  the  next  preparation. 
Unite  the  contents  of  the  absorption  flask  and  absorption 
tube,  and  add  enough  more  6-n  H2SO4  to  just  neutralize  the 
ammonia.  Be  careful  not  to  add  a  drop  too  much  of  the 
H2SO4  because  it  cannot  be  removed  by  evaporation,  — 
it  is  well  to  hold  a  part  of  the  solution  in  reserve  in  case 
the  neutralization  of  the  main  part  is  slightly  overstepped. 
Evaporate  the  solution  to  obtain  solid  ammonium  sulphate, 
following  the  direction  for  Ammonium  Bromide  on  page  202 
of  this  Supplement, 


2IO  SUPPLEMENT 

Questions 

1.  Powdered  aluminum,  unmixed  with  carbon,  does  not 
enter  into  rapid  self-sustaining  action  with  the  elements  of 
the  air.     Why  might  one  expect  that  such  a  reaction  would 
take  place?    Then  how  may  one  explain  its  not  reacting 
according  to  expectation?     Finally  how  can  you  account  for 
the  fact  that  a  small  admixture  of  soot  makes  an  energetic 
reaction  possible? 

2.  What  information  did  you  obtain  in  your   tests  of 
the  gas  passing  unabsorbed  through  the  U-tubes?     What 
probably  is  the  gas  and  how  do  you  account  for  its  forma- 
tion?    (Look  up  Aluminum  Carbide  in  reference  book.) 

3.  Compare    the   hydrolysis   of   aluminum   nitride   with 
that  of  magnesium  nitride  in  No.  5.     Can  you  offer  any 
plausible  explanation  why  the  hydrolysis  of  the  latter  takes 
place  so  much  more  easily?    How  does  the  addition  of  the 
sodium  hydroxide  help  with  the  aluminum  nitride? 

4.  In  what  important  detail  does  the  process  followed 
in  this  laboratory  preparation  differ  from  the  commercial 
process  of  Serpek,  and  why  would  the  former  not  be  feasible 
on  a  commercial  basis?     Discuss  the  possibility  of  the  proc- 
ess outlined  in  5~A  ever  becoming  a  commercial  process  of 
''fixing'''  nitrogen. 

5.  What  is  the  object  of  placing  glass  beads  in  the  second 
U-tube? 

6.  Apply  Question  2  under  No.  5~A  to  aluminum  nitride 
and  aluminum  carbide  as  well  as  to  magnesium  nitride. 

7.  Make  for  this  preparation  a  calculation  similar  to  that 
suggested  by  Question  5  under  No.  5~A. 

II-B.    ALUM  FROM  SODIUM  ALUMINATE 

BY-PRODUCT,    SODA   CRYSTALS,    Na2CO3.IO  H2O 

The  solution  left  in  the  generator  flask  from  the  preced- 
ing preparation  contained  mainly  sodium  meta-aluminate, 


SUPPLEMENT  211 

NaAlO2.  Addition  of  an  acid  to  this  salt  sets  free  the  weak 
meta-aluminic  acid 

H2CO3    +  2  NaA102  -» Na2CO3  +  2  HA1O2 
2  HA1O2  +  2  H2O       -»  2  A1(OH)3 

The  precipitated  aluminum  hydroxide  serves  as  the  starting 
point  in  the  preparation  of  alum.  Evaporation  of  the 
filtrate  yields  soda  crystals  as  a  by-product. 

Read  the  discussion  and  procedure  of  No.  10,  Alum  from 
Cryolite,  and  note  that  we  are  here  starting  with  the  sodium 
aluminate  solution  which  is  the  first  intermediate  product 
in  that  experiment.  Note  further  that  solutions  of  Na3AlO3 
and  NaAlO2  differ  only  by  the  amounts  of  NaOH  they  con- 
tain. 

Apparatus:    large  flask  or  bottle,  2  liters. 

i-hole  stopper  to  fit  mouth  of  same. 
2  pieces  cloth  filter  18  inches  square, 
wooden  stand  12  inches  square  for  filter, 
agate  pail. 
2  8-inch  evaporating  dishes. 

Procedure.  —  To  the  residue  in  the  generating  flask  in 
No.  n-A  add  100  cc.  of  water,  stir,  and  filter  the  solution 
through  paper.  Place  the  clear  solution  in  the  large  flask 
or  bottle  and  nearly  fill  the  latter  with  water.  Through 
the  tight-fitting  stopper  lead  a  gas  delivery  tube  to  the  bottom 
of  the  liquid  and  connect  with  the  automatic  carbon  dioxide 
generator  (Note  12,  page  192,  of  Supplement).  Loosen  the 
stopper  until  all  air  is  expelled,  then  close  the  flask  and  allow 
carbon  dioxide  to  be  absorbed  until  the  solution  is  saturated. 
Tack  one  piece  of  the  cloth  filter  to  the  wooden  frame,  allow- 
ing the  middle  to  hang  4  inches  lower  than  the  edges.  Lay  the 
other  cloth  over  the  first  one.  Collect  the  precipitated 
aluminum  hydroxide  on  this  filter  and  catch  the  filtrate  in  a 


212  SUPPLEMENT 

clean  agate  pail.  Wash  the  precipitate  twice  with  hot  water. 
Transfer  the  A1(OH)3  to  a  dish  and  treat  it  with  three  times 
the  volume  of  6-normal  H2SO4  which  it  took  to  neutralize 
the  ammonia  formed  in  II-A.  Warm  the  mixture  for  5 
minutes.  If  much  undissolved  substance  is  left,  add  5  cc. 
more  of  6-normal  H2SO4,  warm  one  minute,  and  continue 
with  such  additions  until  the  solution  is  nearly  clear.  It  is 
important  to  avoid  adding  more  than  a  trifling  excess  of  acid. 
For  every  100  cc.  of  the  6-normal  acid  used  in  dissolving  the 
A1(OH)3  add  17.4  grams  of  solid  K2SO4.  Warm  until  this  is 
dissolved.  Filter  the  solution  and  bring  it  to  crystallization 
as  directed  in  No.  10. 

Obtain  soda  crystals  from  the  nitrate  in  the  agate  pail. 

n-c.     MODIFICATION  OF   II-B. 

BY-PRODUCT,    GLAUBER   SALT,   Na2SO4.IO   H2O 

Follow  the  same  procedure  as  in  No.  II-B  except  in  using 
H2SO4  instead  of  carbon  dioxide  in  precipitating  aluminic 
acid.  An  excess  of  carbonic  acid  does  no  harm,  but  an 
excess  of  H2SO4  dissolves  an  equivalent  amount  of  A1(OH)3. 
Before  beginning  operations,  submit  to  an  instructor  your 
plan  for  telling  just  how  much  H2SO4  to  add  to  completely 
precipitate  the  A1(OH)3  without  redissolving  any.  Obtain 
crystals  of  alum  and  of  Glauber  salt  as  the  by-product. 

23.   LEAD  NITRATE 

Note  to  Procedure.  —  The  solution  which  is  set  to  crystallize 
should  be  slightly  acid,  —  enough  to  redden  litmus.  If  in- 
sufficient nitric  acid  is  used,  the  excess  of  PbO  dissolves  in 
hot  concentrated  Pb(NO3)2  solution  forming  the  basic  salt 
PbOH.NO3  which  separates  as  a  fine  granular  or  flaky  pre- 
cipitate when  the  solution  cools. 


SUPPLEMENT  213 

35.   CHROMIC  ANHYDRIDE,  CrO3 

Materials:     sodium  dichromate,  Na2Cr2O7.2H2O,  100  grams. 

concentrated  sulphuric  acid,  400  cc. 
Apparatus:   8-inch  evaporating  dish. 

glass  plate  to  cover  the  8-inch  dish. 

glass  marble. 

unglazed  porcelain  plate. 

glass-stoppered  sample  bottle. 

Procedure.  —  Dissolve  the  100  grams  of  sodium  dichromate 
in  250  cc.  of  water  and  filter  from  any  sediment.  Add 
rather  slowly  with  constant  stirring  about  half  of  the  con- 
centrated sulphuric  acid  until  a  slight  permanent  precipitate 
of  CrO3  is  formed.  Let  the  mixture  cool  for  half  an  hour 
or  longer,  then  add  slowly,  while  stirring,  the  rest  of  the  sul- 
phuric acid.  Let  the  mixture  stand  over  night  covered  with 
a  glass  plate  in  order  that  the  crystal  meal  may  become 
somewhat  coarser.  In  such  a  crystal  meal  standing  in  its 
saturated  solution,  the  smaller  grains  dissolve  and  their 
material  deposits  out  on  the  larger  crystals.  But  even 
now  the  crystal  meal  will  be  rather  fine  and  it  will  at 
first  run  through  the  filter;  if,  however,  while  waiting,  the 
mixture  is  heated  with  stirring  to  100°  and  allowed  to  cool 
slowly,  and  this  process  is  repeated  once  or  twice,  a  more 
satisfactory  product  will  be  obtained.  To  collect  the  crys- 
tals, use  a  suction  filter,  but  place  a  small  glass  marble  in 
the  funnel  instead  of  the  usual  plate  and  paper.  If  the  red 
crystals  at  first  run  past  the  sides  of  the  marble,  pour  the 
liquid  in  the  bottle  repeatedly  back  onto  the  filter  until 
finally  the  filtrate  runs  clear  (see  last  sentence  of  Note  3  on 
page  7).  After  draining  the  crystals  completely  and  press- 
ing the  surface  with  a  glass  spatula,  stop  the  suction  and  pour 
15  cc.  of  the  concentrated  nitric  acid  so  as  to  wash  down  the 
sides  of  the  funnel  and  cover  the  surface  of  the  product. 


2 14  SUPPLEMENT 

Stir  up  the  product  with  this  washing  fluid  for  a  depth  of 
about  \  inch.  Suck  dry  and  repeat  the  operation  twice 
with  10  cc.  of  nitric  acid  each  time.  Finally  transfer  the 
product  to  an  unglazed  porcelain  plate,  place  the  latter  on 
an  iron  ring  and  heat  it  by  playing  under  it  the  burner  held 
in  one  hand  while  with  the  other  hand  the  crystals  are  con- 
tinually stirred.  Continue  this  operation,  being  very  careful 
not  to  overheat,  until  nitric  acid  vapors  cease  to  be  given  off. 
Transfer  the  product  at  once  to  a  dry  previously  weighed 
glass-stoppered  bottle. 


364847 


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